Characteristics of group 6 of the main subgroup. Group VI nonmetals

The main subgroup of group VI of the periodic table includes oxygen, sulfur, selenium, tellurium and polonium. The nonmetallic properties of group VI-A elements are less pronounced than those of the halogens. Their valence electrons are ns 2 np 4 .

Oxygen and sulfur are typical nonmetals, with oxygen being one of the most electronegative elements (second only to fluorine). Polonium is a silver-white metal, reminiscent of lead in physical properties, and in electrochemical properties - noble metals. Selenium and tellurium occupy an intermediate position between metals and non-metals; they are semiconductors. In terms of chemical properties, they are closer to non-metals. Oxygen, sulfur, selenium and tellurium are grouped together as “chalcogens,” which translated from Greek means “generating ores.” These elements are found in numerous ores. From oxygen to tellurium, the content of elements on Earth drops sharply. Polonium has no stable isotopes and is found in uranium and thorium ores, as one of the decay products of radioactive uranium.

In their properties, oxygen and sulfur differ sharply from each other, because their electron shells of the previous energy level are constructed differently. Tellurium and polonium have the same structure of the outer energy level (valence layer) and the penultimate energy level, so they are more similar in their properties.

Oxygen is a chemically active non-metal and is the lightest element from the group of chalcogens. The simple substance oxygen under normal conditions is a gas without color, taste or smell, the molecule of which consists of two oxygen atoms (formula O 2), and therefore it is also called dioxygen. Liquid oxygen has a light blue color, and solid oxygen is light blue crystals. There are other allotropic forms of oxygen, for example, ozone - under normal conditions, a blue gas with a specific odor, the molecule of which consists of three oxygen atoms (formula O3 The word oxygen (also called “acid solution” at the beginning of the 19th century) owes its appearance in the Russian language to some extent to M.V. Lomonosov, who introduced the word “acid”, along with other neologisms; thus, the word “oxygen”, in turn, was a tracing of the term “oxygen” (French oxygine), proposed by A. Lavoisier (from the ancient Greek ?oet - “sour” and gennshch - “give birth”), which is translated as “generating acid”, which is related to its original meaning – “acid”, which previously meant substances called oxides according to modern international nomenclature. Oxygen is the most common element in the earth's crust; its share (in various compounds, mainly silicates) accounts for about 47% of the mass of the solid earth's crust. In the atmosphere, the content of free oxygen is 20.95% by volume and 23.10% by mass (about 1015 tons). Currently, in industry, oxygen is obtained from the air. The main industrial method for producing oxygen is cryogenic rectification. Oxygen plants operating on the basis of membrane technology are also well known and successfully used in industry.

Laboratories use industrially produced oxygen, supplied in steel cylinders under a pressure of about 15 MPa.

Small amounts of oxygen can be obtained by heating potassium permanganate KMnO4:

The reaction of catalytic decomposition of hydrogen peroxide H2O2 in the presence of manganese(IV) oxide is also used:

Oxygen can be obtained by the catalytic decomposition of potassium chlorate (Berthollet salt) KClO 3:

Laboratory methods for producing oxygen include the method of electrolysis of aqueous solutions of alkalis, as well as the decomposition of mercury(II) oxide (at t = 100 °C):

In submarines it is usually obtained by the reaction of sodium peroxide and carbon dioxide exhaled by humans:

A strong oxidizing agent, it interacts with almost all elements, forming oxides. Oxidation state?2. As a rule, the oxidation reaction proceeds with the release of heat and accelerates with increasing temperature. Example of reactions occurring at room temperature:

Oxidizes compounds that contain elements with less than the maximum oxidation state:

Oxidizes most organic compounds:

Under certain conditions, it is possible to carry out mild oxidation of an organic compound:

Oxygen reacts directly (under normal conditions, with heating and/or in the presence of catalysts) with all simple substances except Au and inert gases (He, Ne, Ar, Kr, Xe, Rn); reactions with halogens occur under the influence of an electrical discharge or ultraviolet radiation. Oxides of gold and heavy inert gases (Xe, Rn) were obtained indirectly. In all two-element compounds of oxygen with other elements, oxygen plays the role of an oxidizing agent, except for compounds with fluorine.

Oxygen forms peroxides with the oxidation state of the oxygen atom formally equal to?1.

For example, peroxides are produced by the combustion of alkali metals in oxygen:

Some oxides absorb oxygen:

According to the combustion theory developed by A. N. Bach and K. O. Engler, oxidation occurs in two stages with the formation of an intermediate peroxide compound. This intermediate compound can be isolated, for example, when a flame of burning hydrogen is cooled with ice, hydrogen peroxide is formed along with water:

In superoxides, oxygen formally has an oxidation state of ?S, that is, one electron per two oxygen atoms (O ?2 ion). Obtained by reacting peroxides with oxygen at elevated pressure and temperature:

Potassium K, rubidium Rb and cesium Cs react with oxygen to form superoxides:

Inorganic ozonides contain the O?3 ion with the oxidation state of oxygen formally equal to?1/3. Obtained by the action of ozone on alkali metal hydroxides:

Sulfur is an element of the main subgroup of group VI, the third period of the periodic table of chemical elements of D.I. Mendeleev, with atomic number 16. It exhibits non-metallic properties. Denoted by the symbol S (Latin sulfur). In hydrogen and oxygen compounds it is found in various ions and forms many acids and salts. Many sulfur-containing salts are poorly soluble in water. Sulfur is the sixteenth most abundant element in the earth's crust. It is found in a free (native) state and bound form.

The most important natural sulfur minerals: FeS 2 - iron pyrite or pyrite, ZnS - zinc blende or sphalerite (wurtzite), PbS - lead luster or galena, HgS - cinnabar, Sb 2 S 3 - stibnite. In addition, sulfur is present in petroleum, natural coal, natural gases and shale. Sulfur is the sixth most abundant element in natural waters; it is found mainly in the form of sulfate ions and causes the “constant” hardness of fresh water. A vital element for higher organisms, an integral part of many proteins, is concentrated in the hair. The word “sulfur”, known in the Old Russian language since the 15th century, is borrowed from the Old Slavonic “s?ra” - “sulfur, resin”, generally “flammable substance, fat”. The etymology of the word has not been clarified to date, since the original common Slavic name for the substance has been lost and the word has reached the modern Russian language in a distorted form.

According to Vasmer, “sulfur” goes back to lat. sera -- “wax” or lat. serum -- “serum”.

The Latin sulfur (derived from the Hellenized spelling of the etymological sulpur) presumably goes back to the Indo-European root swelp - "to burn". In air, sulfur burns, forming sulfur dioxide - a colorless gas with a pungent odor:

Using spectral analysis, it was established that in fact the process of sulfur oxidation into dioxide is a chain reaction and occurs with the formation of a number of intermediate products: sulfur monoxide S 2 O 2, molecular sulfur S 2, free sulfur atoms S and free radicals of sulfur monoxide SO.

The reducing properties of sulfur are manifested in the reactions of sulfur with other non-metals, however, at room temperature, sulfur reacts only with fluorine.

Molten sulfur reacts with chlorine, and the formation of two lower chlorides (sulfur dichloride and dithiodichloride) is possible.

With an excess of sulfur, various polysulfur dichlorides such as SnCl 2 are also formed.

When heated, sulfur also reacts with phosphorus, forming a mixture of phosphorus sulfides, among which is the higher sulfide P2S5:

In addition, when heated, sulfur reacts with hydrogen, carbon, silicon:

(hydrogen sulfide)
(carbon disulfide)

When heated, sulfur interacts with many metals, often quite violently. Sometimes a mixture of metal and sulfur ignites when ignited. This interaction produces sulfides:

Solutions of alkali metal sulfides react with sulfur to form polysulfides:

Of the complex substances, noteworthy first of all is the reaction of sulfur with molten alkali, in which sulfur is disproportionately similar to chlorine:

The resulting alloy is called liver of sulfur.

Sulfur reacts with concentrated oxidizing acids (HNO 3, H 2 SO 4) only with prolonged heating:

(conc.)
(conc.)

As the temperature increases in sulfur vapor, changes occur in the quantitative molecular composition. The number of atoms in a molecule decreases:

At 800--1400 °C the vapors consist mainly of diatomic sulfur:

And at 1700 °C sulfur becomes atomic:

Sulfur is one of the biogenic elements. Sulfur is part of some amino acids (cysteine, methionine), vitamins (biotin, thiamine), and enzymes. Sulfur is involved in the formation of protein tertiary structure (formation of disulfide bridges). Sulfur is also involved in bacterial photosynthesis (sulfur is part of bacteriochlorophyll, and hydrogen sulfide is a source of hydrogen). Redox reactions of sulfur are a source of energy in chemosynthesis.

A person contains approximately 2 g of sulfur per 1 kg of body weight

Selenium is a chemical element of the 16th group (according to the outdated classification - the main subgroup of group VI), the 4th period in the periodic table, has atomic number 34, denoted by the symbol Se (lat. Selenium), a brittle, shiny, black non-metal (stable allotropic form, unstable form - cinnabar-red). Refers to chalcogens.

The name comes from the Greek. welUnz - Moon. The element is so named due to the fact that in nature it is a satellite of tellurium, which is chemically similar to it (named after the Earth). 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 PbBi2 (S, Se) 3, associated with various sulfides, and sometimes also with cassiterite. Native selenium is occasionally found. Sulfide deposits are of major industrial importance for selenium. The selenium content in sulfides ranges from 7 to 110 g/t. The concentration of selenium in sea water is 4·10?4 mg/l.

Selenium is an analogue of sulfur and exhibits oxidation states? 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.

It is part of the active centers of some proteins in the form of the amino acid selenocysteine. A microelement, but most compounds are quite toxic (hydrogen selenide, selenic and selenous acid) even in moderate concentrations.

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. This role of selenium is constantly growing, demand and prices are growing (hence the shortage of this element).

In modern semiconductor technology, selenides of many elements are used, for example, tin, lead, bismuth, antimony, and lanthanide selenides. The photoelectric and thermoelectric properties of both selenium itself and selenides are especially important.

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 vanadium pentoxide is used in the thermochemical production of hydrogen and oxygen from water (selenium cycle, Lawrence Livermore National Laboratory, Livermore, USA).

The semiconducting properties of selenium in its pure form were widely used in the mid-20th century for the manufacture of rectifiers, especially in military equipment for the following reasons: unlike germanium and silicon, selenium is insensitive to radiation, and, in addition, the selenium rectifier diode has the unique property of self-healing in case of breakdown: the breakdown site evaporates and does not lead to a short circuit, the permissible diode current is slightly reduced, but the product remains functional. The disadvantages of selenium rectifiers include their significant dimensions.

The sixth group of the periodic table of elements consists of two subgroups: the main group - oxygen, sulfur, selenium, tellurium and polonium - and the secondary group - chromium, molybdenum and tungsten. In the main subgroup, a selenium subgroup is distinguished (selenium, tellurium and polonium), a secondary subgroup is called the chromium subgroup. All elements of the main subgroup, except oxygen, can add two electrons, forming electronegative ions.

The elements of the main subgroup have an external electronic

level has six electrons (s2р4). Oxygen atoms have two unpaired electrons and no d-level. Therefore, oxygen exhibits mainly an oxidation state of -2 and only in compounds with fluorine +2. Sulfur, selenium, tellurium and polonium also have six electrons in their outer level (s2p4), but they all have an unfilled d-level, so they can have up to six unpaired electrons and exhibit oxidation states of -2, +4 and + in compounds. 6.

The pattern of changes in the activity of these elements is the same as in the subgroup of halogens: tellurides are most easily oxidized, then selenides and sulfides. Of the oxygen compounds of sulfur, the most stable are sulfur (VI) compounds, and for tellurium - tellurium (IV) compounds. Selenium compounds occupy an intermediate position.

Selenium and tellurium, as well as their compounds with certain metals (indium, thallium, etc.) have semiconductor properties and are widely used in radio electronics. Selenium and tellurium compounds are very toxic. They are used in the glass industry to produce colored (red and brown) glasses.

In elements of the chromium subgroup, the d-level is filled, therefore, at the s-level of their atoms there are one (for chromium and molybdenum) or two (for tungsten) electrons. All of them exhibit a maximum oxidation state of +6, but molybdenum, and especially chromium, are characterized by compounds in which they have a lower oxidation state (+4 for molybdenum and +3 or +2 for chromium). Chromium(III) compounds are very stable and similar to aluminum compounds. All metals of the chromium subgroup are widely used.

Molybdenum was first obtained by K.V. Scheele in 1778. It is used in the production of high-strength and toughness steels used for the manufacture of weapon barrels, armor, shafts, etc. Due to the ability to evaporate at high temperatures, it is of little use for the manufacture of filaments , but has a good ability to fuse with glass, so it is used to make tungsten filament holders in incandescent lamps.

Tungsten was also discovered by K.V. Scheele in 178! d. It is used to produce special steels. The addition of tungsten to steel increases its hardness, elasticity and strength. Together with chromium, tungsten gives steel the ability to maintain hardness at very high temperatures, which is why such steels are used to make cutters for high-speed lathes. Pure tungsten has the highest melting point among metals (3370 °C), therefore it is used to make filaments in incandescent lamps. Tungsten carbide is characterized by very high hardness and heat resistance and is the main component of refractory alloys.

72. Oxygen

Oxygen was discovered by the Swedish chemist K.V. Scheele in 1769-1770. and the English chemist D. J. Priestley in 1774

Being in nature. Oxygen is the most abundant element in nature. Its content in the earth's crust is 47.00% by weight. In a free state, it is found in the atmosphere (about 23% by weight), is part of water (88.9%), all oxides that make up the earth’s crust, oxygen-containing salts, as well as many organic substances of plant and animal origin.

HEDGEHOG (anabasis), a genus of perennial herbs or subshrubs of the goosefoot family. OK. 30 species, in the Center. Asia, in the south of Europe, in the North. Africa, but mainly in Wed. Asia. Some species are pasture food for camels and sheep. Sometimes some species of the cereal family (paizu, chicken millet) are also called barnyard grass.

NILSBOHRIUM (lat. Nielsbohrium), Ns, artificially obtained radioactive chemical element V gr. periodic table, atomic number 105. The most stable isotope is 262Ns (half-life 40 s). Obtained in 1970 in the USSR and the USA. Named at the suggestion of Soviet physicists after Niels Bohr; American scientists proposed the name "ganium" in honor of O. Gan. The name has not been finalized.

KOKOV Valery Mukhamedovich (b. 1941), Russian statesman, president of the Kabardino-Balkarian Republic (1993). In 1990-91, Chairman of the Supreme Council of the Kabardino-Balkarian Autonomous Soviet Socialist Republic, in 1991-92, First Deputy Chairman of the Council of Ministers of the Kabardino-Balkarian Republic. In 1993-95, deputy of the Federation Council of the Federal Assembly of the Russian Federation, since 1995 member of the Federation Council. In 1997 he was elected president for a second term.

Oxygen, sulfur, selenium, tellurium and polonium form the main subgroup of the sixth group of the periodic table and are p-elements. Their atoms have six electrons in the outer electronic level, and the overall electronic configuration of the outer electronic layer can be expressed by the formula: ns2np4. Electronic formulas of atoms and some physical constants are given in the table.

electron configuration of an atom	average atomic mass	apparent radius of a neutral atom, A	electron affinity, eV	relative electro negativity	apparent ion radius

The following conclusions follow from the table data:

1. The apparent radii of neutral atoms and negative ions increase correctly with increasing atomic number of the element.

2. The value of relative electronegativity decreases with increasing apparent radii of neutral atoms. Consequently, from oxygen to polonium, the oxidizing properties weaken and the reducing properties of neutral atoms increase. The strongest oxidizing agent among these elements is oxygen:

O – S – Se – Te – Po

Strengthening oxidative properties

3. With an increase in the serial numbers of elements, a gradual weakening of non-metallic properties and an increase in metallic properties are observed.

The distribution of valence electrons for p-elements of the sixth group over atomic orbitals has the following form:

for oxygen

for sulfur, selenium, tellurium and polonium

The presence of six electrons on the outer quantum layer characterizes the ability of the elements under consideration to exhibit a negative oxidation state 2–. All elements are capable of forming negatively charged ions with a charge of 2–. The tendency to form negatively charged E-2 ions weakens from oxygen to polonium.

The oxygen atom lacks a d-sublevel. Therefore, due to the presence of two unpaired p-electrons, the oxygen atom can form two chemical bonds with atoms of other elements. From this it is clear that compounds formed by oxygen with monovalent elements have the formula E2O. In addition, an oxygen atom can form a bond via a donor-acceptor mechanism.

The oxygen atom can act as a donor - due to the unshared pair of electrons it has, for example, during the formation of the hydronium ion (H2O + H+ = H3O+) and as an acceptor - due to the free orbital that appears upon excitation by pairing of two unpaired ones electrons (which is observed, for example, in the nitric acid molecule // O

N – O – N.

Depending on the nature of the atom with which oxygen interacts, the degree of its oxidation can be different:

2(H2O); -1(H2O2); 0(O2); +1(O2F2); +2(OF2).

The atoms of sulfur, selenium, tellurium and polonium have a free d-sublevel. When these atoms are excited, their electrons can move to vacant d-orbitals and therefore these elements exhibit the following oxidation states: -2, +2, +4, +6.

Simple substances.

A feature of this group is the polyatomic nature of the molecules of simple substances En, where 2 ≤ n ≤ ∞.

	oxygen
composition of molecules		S8 (room tº)	Se8; Se∞ (room tº)	Te∞(room tº) Te2(>1400ºC)
allotropic modifications	O2(oxygen)	rhombic (below 95.6ºC) monoclinic (above 95.6ºC) amorphous (plastic)	red (crystalline metal (gray) amorphous	metal amorphous	α - modified β-modified
ρ, density g/cm3			4.82 (metal)	6.25 (metal)
melting temperature ºC
boiling point ºC
element prevalence	Clark - 49% Lithosphere-47.3% air -23.1%

From the data presented in the table, the following conclusions can be drawn:

1. Molecules of simple substances formed by atoms of p-elements of group VI are polyatomic.

2. All elements are characterized by the presence of allotropic modifications.

3. Boiling and melting points (except for polonium), their densities increase with increasing serial number.

Sulfur was known before 5000 BC. e.

Oxygen was discovered four times: in 1772 by Scheele (HgO), in 1774 by Pierre Bayen, Priestley (calcined Pb3O4 - got PbO and O2), Lavoisier gave the name to oxygen and determined that it is part of the air.

Tellurium (earth) – 1798 Klaproth.

Selenium (moon) – 1817 Berzelius found in sludge

Polonium - 1898 discovered by Marie Curie - Skladowska and named after her homeland (Poland).

Oxygen

Oxygen is the most abundant element on Earth (58.0 mole fractions). Its great chemical activity and quantitative predominance largely determine the forms of existence of all other elements on Earth. The most common natural oxygen compounds are H2O, SiO2, silicates and aluminosilicates. In the air, oxygen is in a free state and accounts for 20.99% (vol.). In the upper layers of the atmosphere, oxygen is found in the form of a gas - ozone (O3). The ozone layer traps hard solar radiation, which, with prolonged exposure to living organisms, is fatal to them.

Natural oxygen consists of three stable isotopes: 16O (99.795%), 17O (0.037%) and 18O (0.204%). In addition, three radioactive isotopes were obtained, the lifetime of which is negligible.

In terms of relative electronegativity (REO = 3.5), oxygen is second only to fluorine. Oxygen forms compounds with almost all elements, excluding helium, neon and argon. In compounds with other elements, in addition to the oxidation states already mentioned (-2, -1, +1, +2), oxygen exhibits an oxidation state of +4 in ozone.

For oxygen, two allotropic modifications are known: 1) O2 – oxygen; 2) O3 – ozone.

The most stable diatomic molecule is oxygen (O2). The bond order in this molecule is 2. From the energy diagram it follows that oxygen is a paramagnetic substance (the molecule has two unpaired electrons). This position is fully confirmed by experience. The dissociation energy of the O2 molecule is 494 kJ/mol, which indicates its sufficient stability. The chemical activity of the oxygen molecule is explained by the presence of unpaired electrons in antibonding π orbitals. Under normal conditions, O2 is a colorless gas. Liquid oxygen is blue in color. Crystals of solid oxygen are colored light blue and look like snow. Oxygen is slightly heavier than air (dair = 1.105). Oxygen dissolves in water in very small quantities. In every solid state, oxygen is attracted by a magnet.

Obtaining oxygen

In industry, oxygen is obtained from liquid air, by electrolysis of water, as a by-product in the production of high-purity hydrogen.

In the laboratory, oxygen is obtained from the thermal decomposition of oxygen-rich compounds (KМnO4, KСlO3, KNO3, etc.).

For example: 2КMn+7O4-2 tº→ К2Мn+6О4 + Mn+4О2 + О20

Such reactions belong to intramolecular oxidation-reduction reactions.

Chemical properties

In terms of reactivity, O2 is second only to halogens. Its chemical activity increases with increasing temperature. O2 interacts with almost all chemical elements, with the exception of halogens, noble gases and noble metals (silver, gold, platinum). Sometimes the interaction is prevented by an oxide film on the surface of the oxidized substance.

The rate of oxidation reactions depends on the nature of the substance being oxidized, temperature, catalyst, etc. Most oxidation reactions are exothermic, e.g.

C + O2 → CO2 ΔΗ = -382.5 kJ/mol

2H2 + O2 → 2H2O ΔΗ = -571.7 kJ/mol

Application of oxygen

The bulk of the oxygen produced by industry is consumed in ferrous metallurgy to intensify the smelting of iron and steel. Oxygen is widely used in petrochemical
industry" href="/text/category/himicheskaya_i_neftehimicheskaya_promishlennostmz/" rel="bookmark">chemical industry for the production of sulfuric and nitric acids, lubricating oils, etc. Mixed with acetylene, O2 is used for welding and cutting metals (flame temperature approx. 3200ºC). Liquid oxygen is used in rockets and mining.

Ozone

Ozone (O3) is the second allotropic modification of oxygen. It is a blue gas with a pungent odor (bp -112ºC, mp -193ºC). Liquid ozone is a dark blue liquid. Solid ozone is black. Ozone is highly toxic and explosive. The formation of ozone molecules is accompanied by the absorption of energy:

https://pandia.ru/text/78/050/images/image014_50.gif" width="50" height="51 src=">O

https://pandia.ru/text/78/050/images/image017_44.gif" width="38" height="38"> 126 Ǻ 116.5º

Ozone is produced by the action of a quiet electrical discharge on oxygen. A small amount of ozone is formed in processes accompanied by the release of atomic oxygen (radiolysis of water, decomposition of peroxides, etc.). Under natural conditions, ozone is formed from atmospheric oxygen during lightning discharges and under the influence of ultraviolet rays from the sun. The maximum ozone concentration is formed at an altitude of ≈ 25 km. The “ozone belt” plays a vital role in ensuring life on Earth, as it blocks ultraviolet radiation harmful to living organisms and absorbs infrared radiation from the Earth, preventing its cooling.

Ozone is a more active oxidizing agent than oxygen. For example, already under normal conditions it oxidizes many metals and other substances

2Ag + O3 → Ag2O + O2

PbS + 4O3 → PbSO4 + 4O2

Reactions involving it usually produce oxygen. Ozone reacts with many substances under conditions where oxygen remains inert. Thus, the reaction O3 + 2KI + H2O = I2 + 2KOH + O2 occurs quantitatively and can be used for the quantitative determination of ozone.

In addition, reactions are known in which the ozone molecule participates with all three oxygen atoms, for example KI + O3 → KIO3.

3SnCl2 + O3 + 6HCl = 3 SnCl4 + 3H2O.

The use of ozone is due to its oxidizing properties. It is used as a disinfectant and bactericide, for water purification, in the food industry, etc.

Peroxides

Peroxides are oxygen compounds in which the oxygen atoms are directly bonded to each other. Thus, in the structure of peroxides there is a group –O–O–, it is called peroxide ion.

Peroxide and superoxide ions are produced by combining electrons with an O2 molecule

O20+e → O2- - superoxide

O20+2e → 2O2- -peroxide

O2-pair O2-pair O22-dia -

decrease in stability

Compounds containing superoxide ion (O2-) are called superoxides, for example, KO2. The presence of an unpaired electron in them determines their paramagnetism. The peroxide ion (O2-2) has no unpaired electrons and therefore this ion is diamagnetic. In peroxides, oxygen atoms are connected to each other by one two-electron bond. The formation of peroxides is typical for active metals (alkali, alkaline earth). Hydrogen peroxide (H2O2) is of most practical importance.

The H2O2 molecule is polar (μ=0.70∙10-29 C∙m.) The presence of hydrogen bonds causes the high viscosity of hydrogen peroxide. Due to the association of molecules, H2O2 under normal conditions is a liquid (tmelt = -0.410C, tbp = 1500C). Hydrogen peroxide easily decomposes into atomic hydrogen and oxygen, H2O2=t H2+O2

soluble in water, aqueous solution of H2O2 is a weak acid. Dissociation constant Kg(H2O2)=2.24∙10-12

Hydrogen peroxide can be obtained using the general method for producing weak acids (displacing a weak acid from its salt with a stronger acid)

BaO2+H2SO4=H2O2+BaSO4↓

In industry, hydrogen peroxide is produced by electrochemical oxidation of sulfuric acid at low temperature on a platinum anode.

H2 O2 in OVR

Oxygen in hydrogen peroxide is assigned an oxidation state of –1 (this oxidation state is intermediate to oxygen). Therefore, it can be both an oxidizing agent and a reducing agent. When H2O2 is reduced, water or OH- is formed, e.g.

2KJ+H2O2+H2SO4=J2+K2SO4+2H2O

PbS+H2O2=PbSO4+H2O

2K3+3H2O2=2K2CrO4+2KOH+8H2O

In these cases, the process occurs: H2O2 oxidizer

When interacting with strong oxidizing agents, hydrogen peroxide exhibits the properties of a reducing agent.

5H2O2+2KMnO4+3H2SO4→5O20+2MnSO4+K2SO4+8H2O

This reaction is used in chemical analysis to quantify the content of hydrogen peroxide in a solution.

Application:

3% solution in medicine as an antiseptic,

6%-12% solution – for hair bleaching,

more than 30% conc. in the chemical industry.

Sulfur

General characteristics of sulfur. Unlike oxygen, sulfur has vacant 3d orbitals in its outer quantum layer.

Sulfur can have the following oxidation states:

2 (H2S, H2S2O3 and sulfides, sodium thiosulfate Na2S2O3 5H2O, where one sulfur atom has an oxidation state of –2 and the other +6.;

2 (S2Cl2, 3SO→SO2+S2O)

4 (SO2, H2S+4O3, its salts);

6 (SO3, H2SO4, its salts, H2S2O7 pyrosulfuric acid)

H2SO5-peroxomonosulfuric acid

H2S2O8-peroxodisulfuric acid

Sulfur is a typical non-metal (oeo = 2.5), it is chemically active and directly combines with almost all elements, with the exception of nitrogen, iodine, gold, platinum, and noble gases. It occurs in nature both in a free state (native sulfur) and in the form of various compounds.

Native sulfur is rare; the most common minerals are sulfide (FeS2, CuS, ZnS, Sb2S3, AgS) and sulfate compounds (CaSO4 2H2O, BaSO4, MgSO4 7H2O, Na2SO4 10H2O), SO2, H2S - contain volcanic gases. In addition, sulfur is part of plant and animal proteins and compounds found in oil. In all solid and liquid states, sulfur is diamagnetic.

Simple substances

Sulfur exists in several allotropic modifications. At room temperature, yellow orthorhombic sulfur (α-S) is stable and consists of very small crystals. Large crystals of this form can be obtained by slow crystallization of sulfur from a solution of sulfur in carbon disulfide. They turn out correctly cut and transparent.

The second allotropic modification is monoclinic sulfur (β-S) needle-shaped crystals.

Allotropic modifications of α- and β-sulfur consist of S8 molecules, which have a cyclic “toothed” structure.

In rhombic sulfur, the rings are 3.3 Ǻ apart. They are interconnected by van der Waals forces. This modification does not conduct heat or electricity.

The difference in the physical properties of orthorhombic and monoclinic sulfur is not due to different molecular compositions (both are composed of S8), but to different crystal structures.

There are other allotropic modifications of sulfur that are formed when temperature changes. Changing pressure also gives different allotropic forms.

With increasing temperature, sulfur changes its color as the chain length decreases:

600ºC 900ºC 1500ºC

orange red yellow

The most stable modification is the rhombic modification; all other modifications spontaneously transform into it.

Sulfur is highly soluble in organic solvents, especially in carbon disulfide and benzene (34%, t = 25ºC).

Chemical properties of sulfur

Sulfur is a very active element. When interacting with stronger oxidizing agents (O2, CI2, etc.), it can give up its electrons, that is, be a reducing agent:

S + Cl2 ↔ SCl2 (S2Br2, S2Cl2)

S0 - 4ē → S+4

2O0 + 4ē → 2O-2

P4 + xS ↔P4Sx x ~ 3, x~ 7

When melted or heated, sulfur reacts with almost all metals to form non-stoichiometric compounds (exhibits the properties of an oxidizing agent).

Hg0 + S0 = Hg+2S-2

Sulfur reacts with most metals when heated, and with mercury at room temperature. Therefore, spilled mercury is covered with sulfur in order to disinfect the room from mercury vapor.

Sulfur is also prone to disproportionation reactions. For example, when powdered sulfur is boiled in an alkali solution, the reaction occurs

S0 + 2S0 +6NaOH = Na2+4SO3 + 2Na2S-2 + 3H2O

Sulfur reacts with acids

S + 2H2SO4(conc) = 3SO2 + 2H2O

S + 6HNO3(conc) = H2SO4 + 6NO2 +2H2O

Obtaining sulfur

In industry, sulfur is obtained by separating it from waste rock using hot water under high pressure. Sulfur is obtained by chemical methods as follows:

1. From waste gases of metallurgical and coke ovens

2H2S + SO2 → 3S + 2H2O

2. From natural sulfates by calcining them with coal (the process occurs in several stages)

CaSO4 + 4C = 4CO + CaS

CaS + HOH + CO2 = CaCO3 + H2S

Hydrogen sulfide is burned:

2H2S + O2 = 2S↓ + 2H2O

Application of sulfur

Sulfur is used for the production of organic sulfur dyes (CS2), carbon disulfide, in the production of artificial fibers, explosives, and in processes for the production of sulfuric acid.

Sulfites and hydrosulfites are used as reducing agents. Calcium hydrosulfite Ca(HSO3)2 is used in the production of cellulose.

Sulfur compounds with oxidation state +6

Sulfur exhibits the +6 oxidation state in compounds with oxygen and halogens. The most typical compound is sulfur trioxide SO3. In the SO3 molecule, sulfur is in a state of sp2 hybridization. The molecule is a flat triangle. ∟O-S-O = 120º; the molecule is nonpolar (μ = 0).

In the SO3 molecule there are 3π bonds per 3σ bond. The molecule is strong, but less so than SO2. The SO3 molecule polymerizes easily. Under normal conditions, SO3 is a liquid (bp 44.8˚C), which solidifies into a transparent mass (mp 16.8 ºC). SO3 is a typical acidic oxide and reacts vigorously with basic oxides. SO3 reacts vigorously with water to form sulfuric acid and release a large amount of heat.

SO3 + H2O = H2SO4, ΔH = -87.8 kJ

SO3 is used as a sulfonating agent in organic synthesis, as a dehydrating agent in the production of HNO3, for the preparation of oleum, etc.

Sulfuric acid

H2SO4 is a strong dibasic acid. It is a derivative of tetraoxosulfate (VI) – 2- ion. In the 2- ion, sulfur is in a state of sp3 hybridization (4 σ bonds + 2π bonds). Ion 2- has the shape of a regular tetrahedron. The S-O bond length is 1.49 Å. This connection is strong.

H2SO4 is highly soluble in water, and a large amount of heat is released due to the formation of hydrates of the composition H2SO4·H2O, H2SO4·2H2O, H2SO4·4H2O. In this regard, when preparing H2SO4 solutions, the acid should be carefully poured into water in a thin stream, and not vice versa. Concentrated H2SO4 energetically attracts moisture and is therefore used for drying gases. This also explains the charring of many organic compounds (carbohydrates).

C12H22O11 + H2SO4 = 12C + H2SO4∙11H2O

Concentrated sulfuric acid can absorb SO3 in large quantities, forming pyrosulfuric acid H2S2O7. Such solutions are called oleum. In oleum there is an equilibrium of H2SO4 + SO3 H2S2O7.

Sulfuric acid forms two types of salts: medium (sulfates) Me2+1SO4 and acidic (hydrosulfates) Me+1HSO4. Most sulfates are highly soluble in water. Sparely soluble sulfates include Ba(II), Ca(II), Sr(II), Pb(II).

The formation of a sparingly soluble white precipitate of BaSO4 is an analytical reaction to the sulfate ion.

SO42- + Ba2+ = BaSO4 (white crystalline precipitate)

BaSO4 is insoluble in hydrochloric acid.

Some sulfates containing water of crystallization are called vitriol. The latter include CuSO4·5H2O (copper sulfate - blue), FeSO4·7H2O (iron sulfate - green).

Among the salts of sulfuric acid, the crystalline hydrates of its double salts are interesting - alum of the general formula Me2+1SO4 Me2(SO4)3 24H2O, where Me+1(Na, K,NH4, etc.), Me+3(Al, Cr, Fe, Co, etc.).

The most widely known are: aluminum-potassium alum KAl(SO4)2·12H2O, chromium-potassium alum KCr(SO4)2·12H2O, iron-ammonium (NH4)2·Fe2(SO4)3·24H2O. Alum is used as a tanning agent in the leather industry, as a mordant when dyeing fabrics, in medicine, etc.

FunctionS (VI) in redox reactions

The +6 oxidation state is the highest for sulfur, and therefore S+6 functions in redox reactions only as an oxidizing agent.

The oxidizing properties of sulfur (+6) appear only in concentrated sulfuric acid. In dilute sulfuric acid, the oxidizing agent is the H+ proton. Concentrated sulfuric acid is a fairly strong oxidizing agent. It oxidizes non-metals (C, S, P) to higher oxides.

S+2H2SO4 = 3SO2 + 2H2O

C + 2H2SO4 = CO2 + 2SO2 +2H2O

HBr and HI are reduced by sulfuric acid to free halogens

8HI + H2SO4 = 4I2 + H2S + 4H2O

2HBr +H2SO4 = Br2 + SO2 +2H2O

Concentrated sulfuric acid oxidizes many metals (except gold and platinum). Concentrated iron sulfuric acid is passivated and therefore can be transported in steel cylinders. The products of the reduction of concentrated sulfuric acid can be various sulfur compounds. Sequential series of sulfuric acid reduction

H2S+6O4→S+2O2→S0→H2S-2

The nature of the reduction products will depend on the activity of the metal: the more active the metal, the deeper the reduction of sulfur (VI).

5H2SO4conc + 4Mg = 4MgSO4 + H2S+ 4H2O

2H2SO4conc + Cu = CuSO4 + SO2 + 2H2O

Zn + 2H2SO4conc = ZnSO4 + SO2 + 2H2O

When dilute sulfuric acid acts on metals, the reduction product is H2 and only metals up to hydrogen in the electrochemical series dissolve in dilute sulfuric acid.

H2SO4dissolved + Zn = ZnSO4 + H2

3Zn + 4H2SO4dil = S↓ + 3ZnSO4 + 4H2O

4Zn + 5H2SO4 very diluted = H2S + 4ZnSO4 + 4H2O

Preparation of sulfuric acid

The essence of the industrial method for producing sulfuric acid is the oxidation of sulfur dioxide SO2 to sulfur trioxide SO3 and the conversion of the latter into sulfuric acid. The production scheme can be presented as follows:

FeS2 SO2 SO3 H2SO4

This process is carried out in two ways: contact and nitrous. In the contact method for producing sulfuric acid, vanadium anhydride V2O5 with the addition of K2SO4 or PbSO4 is used as a catalyst for the oxidation of SO2. In the nitrous method of producing sulfuric acid, the catalyst that accelerates the oxidation of SO2 to SO3 is nitric oxide NO.

Application of sulfuric acid

Sulfuric acid is one of the most important products of the basic chemical industry. Most chemical compounds are obtained with the direct or indirect participation of sulfuric acid. Sulfuric acid is widely used in the production of mineral fertilizers.

It is used to obtain many mineral acids and salts, and is used in organic synthesis, in the production of explosives, dyes, in textiles, leather and other industries.

Peroxosulfuric acids are sulfur oxygen acids characterized by the presence of a peroxo group – O-O. Two sulfur peroxoacids are well known: peroxomonosulfur H2SO5 and peroxodisulfur H2S2O8.

Peroxomonosulfuric acid (Caro acid) H2SO5 is the peroxide form of sulfuric acid

H – O – O – S – O – H

H2SO5 is one of the strong monobasic acids. Like hydrogen peroxide, it is unstable and is a very strong oxidizing agent.

2KI + H2SO5 = K2SO4 + I2 + H2O

H2SO5 is obtained as an oxidizing agent in organic synthesis. Peroxydisulfuric acid H2S2O8 has the structure

H – O – S – O – O – S – O – H

It also belongs to hydrogen peroxide derivatives and is a very strong oxidizing agent (can oxidize Cr+3 → Cr+6, Mn+2 → Mn+7, 2I - → I0)

2KI + H2S2O8 = 2KHSO4 + I2

H2SO5 and H2S2O8 hydrolyze to form hydrogen peroxide and therefore are used in the industrial production of H2O2 solutions

H2S2O8 + 2H2O = 2H2SO4 + H2O2

H2SO5 + H2O = H2SO4 + H2O2

Sulfur thioacids

Thioacids are derivatives of oxygen acids in which some or all of the oxygen atoms are replaced by sulfur. Salts of thioacids are called thiosalts. An example of thioacids is thiosulfuric acid H2S2O3, which is a derivative of sulfuric acid in which one oxygen atom is replaced by a sulfur atom. Its structural formula has the form

Na2SO3S-2 + 4Cl2 + 5H2O = 2H2SO4 + 6HCl + 2NaCl

Na2S2O3 + Cl2 + H2O = S↓ + Na2SO4 + 2HCl.

When thiosulfate interacts with weak oxidizing agents (I2, Fe3+ and others), the tetrathionate ion S4O62- is formed. The reaction between sodium thiosulfate and iron(III) salts is used to detect thiosulfate ions. The reaction proceeds as follows

2Na2S2O3 + I2 = Na2S4O6 + 2NaI

2FeCl3 + 2Na2S2O3 = 2FCl2 + Na2S4O6 + 2NaCl

When this reaction occurs, an intermediate compound is formed, colored dark purple - Cl. This is an unstable Fe3+ complex, which quickly decomposes by an intramolecular oxidation-reduction reaction according to the scheme

2+ = 2Fe2+ + S4O62-

In this case, the color disappears.

In addition, H2S2O3 is characterized by reactions that proceed through the mechanism of intramolecular oxidation-reduction

H2S2O3 = H2+4SO4 + S0

This explains the instability of thiosulfuric acid. Sodium thiosulfate is used in photography (fixative), in the textile industry, and medicine.

Connecting their character

O3 is an allotropic type of oxygen change, obtained from oxygen using a so-called ozonizer 3O2 = 2O3 - a blue gas.

Complete the equations.

Water - H2O - weak electrolyte, pure water, colorless, odorless liquid, boiling point - 1000, freezing point - 00C, density 1 g/ml. It occurs in three states of aggregation. Water is purified by sublimation in a distiller; the resulting water is called distilled water.

Sulfur oxide (IV) - SO2 - with a pungent suffocating odor, colorless gas.

When dissolved in water, sulfurous acid is formed: SO2 + H2O= H2SO3

VI-SO3-sulfur dioxide is a colorless liquid. When interacting with water, sulfuric acid is formed.

Sulfuric acid

H2 SO4 is a colorless, highly soluble liquid in water.

View document contents
“General characteristics of non-metals. Elements of group VI A »

Lesson Plan No. 15

Date Subject chemistry group

FULL NAME. teacher: Kayyrbekova I.A.

I. Lesson topic: General characteristics of non-metals. Elements of group VI A. Oxygen. Water. Sulfur. Hydrogen sulfide. Sulfuric acid and sulfates.

Type of lesson: lesson learning new knowledge

Target:. Characteristics of chemical elements of group V I A. Be able to describe and prove the chemical properties of sulfur using an example. To acquaint students with the structure and general properties of non-metals, based on their position in the periodic system of atomic structure. Know some ways to obtain non-metals. Be able to give a general description of nonmetals based on their position in the PS and the structure of their atoms.

Tasks:

A) Educational: repeat and systematize students’ knowledge about the properties of chemical elements of the sixth group, about the structure of the atom and the use of compounds;

consolidate the ability to solve calculation problems using reaction equations;

Ә) Educational: conduct environmental education in a chemistry lesson.

B) Developmental: continue the development of logical thinking, the ability to use theoretical knowledge in new situations;

strengthen the skills of comparing, contrasting, analyzing;

II. Expected results:

A) Students should know: Characteristics of chemical elements of group V I A.

Ә) Students should be able to: Be able to describe and prove, using an example, the chemical properties of sulfur

b) students must master: To acquaint students with the structure and general properties of non-metals, based on their position in the periodic system of atomic structure

III. Method and techniques for each stage of the lesson: problem, search, laboratory work, independent work of students.

IV. Facilities: interactive board

During the classes

I. Organizational part Check student attendance. Familiarization with the purpose and objectives of today's lesson. Setting the lesson goal.

II. Updating basic knowledge:

A) Check your notebooks

D/z 153 page No. 2

158 page No. 6 exercise

Task Calculate the volume of gas that will be released when 19.5 g of potassium reacts with phosphoric acid -

B) Independent work

ІІІ. Explanations of new materialand consolidation of new material

Plan:

General characteristics of non-metals - elements of group VI A.

These include oxygen, sulfur, selenium, tellurium, and polonium. Polonium is a radioactive metal, and the rest are ore-forming chalcogens. Of these, oxygen and sulfur are especially important. At the last energy level they have 6 electrons, the highest oxidation state is +6, +2 is constant for oxygen, the lowest is 2 for sulfur. General oxide formula RO 3 and hydrogen compounds with general formula RH 2

characteristic	oxygen		sulfur
Position in P.S. - 1 point	2nd small period, element VI A of group		element VI A group, 3 small periods.
Atomic structure-1 point	O (8p + ;8n 0)8e - 1s 2 2s 2 2p 4		S(16p + ;16n 0)16e - 1s 2 2s 2 2p 6 3s 2 3p 4
Being in nature	It is found in the form of a compound and in a free form. It is part of atmospheric air and is formed as a result of photosynthesis.		found in compound and free form.
receiving	In the laboratory: 2KMnO 4 = K 2 MnO 4 +MnO 2 +O 2 2KClO 3 =2KCl+3O 2 In industry, liquid oxygen is obtained from air.
physical properties	Colorless and odorless gas		Sulfur is a yellow, solid, crystalline substance. 3 types of allotropic modifications: rhombus, monoclinic and plastic.
Chemical properties	When heated it reacts: with carbon with phosphorus - they form oxides. With hydrogen At t= 1500 0 C with nitrogen Does not react directly with halogens. At normal temperatures it reacts with active metals. Reacts with low-reactive metals when heated Complex substances burn and oxides of these elements are formed. The task is to give examples and complete the reaction equationmax-8 points		With simple substances: oxygen, halogen, metals, alkali. Assignment to complete reaction equations 4 points
Connecting their character
O 3 is an allotropic type of change in oxygen; it is obtained from oxygen using a so-called ozonizer 3O 2 = 2O 3 - a blue gas.		Hydrogen sulfide is a poisonous, colorless gas with a rotten egg odor. Subject to dissociation, burns in air, with acids. Complete the equations. Obtained by reacting iron sulfide with dissolved hydrochloric acid.
Water - H 2 O - weak electrolyte, pure water, colorless, odorless liquid, boiling point - 100 0, freezing point - 0 0 C, density is 1 g / ml. It occurs in three states of aggregation. Water is purified by sublimation in a distiller; the resulting water is called distilled water. At ordinary temperatures, it reacts with active metals, basic and acidic oxides. And also with some salts, which results in the formation of crystalline hydrates.		Sulfur oxide (IV) - SO 2 - with a sharp suffocating odor, colorless gas. When dissolved in water, sulfurous acid is formed: SO 2 + H 2 O= H 2 SO 3 Obtained by burning sulfur in air or burning pyrite. VI- SO 3 - sulfur dioxide is a colorless liquid. When interacting with water, sulfuric acid is formed.
Sulfuric acid
H 2 SO 4 is a colorless, highly soluble liquid in water. Used to produce hydrochloric, hydrofluoric, nitric, and phosphoric acid. Dissolved acid reacts with metals up to N. Concentrated acid - reacts with metals, non-metals

Fastenings: 177 page No. 12upr

D/z 153 page No. 2, 158 page No. 6 exr. Abstract The role of oxygen in nature. Application of oxygen

Oxygen in its compounds usually exhibits a valence of two. But in principle, it can also be four-valent, since oxygen has 2 unpaired electrons and 2 lone electron pairs on the outer layer. But since the oxygen atom is small in size, the maximum valency of oxygen is three, since only three hydrogen atoms can fit around it.

Files: 1 file

General characteristics of elements of VI A subgroup

Since the atoms of group VI-A elements contain six electrons on the outer layer, they tend to fill the outer energy level with electrons and are characterized by the formation of E2- anions. The atoms of the elements under consideration (except for polonium) are not inclined to form cations. Oxygen and sulfur are typical nonmetals, with oxygen being one of the most electronegative elements (second only to fluorine). Polonium is a silvery-white metal, reminiscent of lead in physical properties, and noble metals in electrochemical properties. Selenium and tellurium occupy an intermediate position between metals and non-metals; they are semiconductors. In terms of chemical properties, they are closer to non-metals. Oxygen, sulfur, selenium and tellurium are grouped together as “chalcogens,” which translated from Greek means “generating ores.” These elements are found in numerous ores. From oxygen to tellurium, the content of elements on Earth drops sharply. Polonium has no stable isotopes and is found in uranium and thorium ores, as one of the decay products of radioactive uranium.

Oxygen and its compounds

Properties of oxygen. Oxygen O2 is a colorless, odorless and tasteless gas. Poorly soluble in water: at 20°C, about 3 volumes of oxygen dissolve in 100 volumes of water. Liquid oxygen has a light blue color and is attracted by a magnet because its molecules are paramagnetic and have two unpaired electrons. The bond energy in the O2 molecule is 493 kJ/mol, the bond length is 0.1207 nm, and the bond order in the molecule is two. In nature, oxygen exists in the form of three isotopes 16O, 17O, 18O and in the form of two allotropic modifications of oxygen O2 and ozone O3. The air contains about 21% free oxygen.

Obtaining oxygen. In the laboratory, oxygen is obtained by the decomposition of compounds rich in oxygen: a) 2 KClO3 = 2 KCl + 3 O2 (catalyst - MnO2) b) 2 KMnO4 = O2 + K2MnO4 + MnO2 c) H2O2 = 2 H2O + O2 (catalyst - MnO2) d) electrolysis of aqueous solutions of oxygen-containing acids and alkalis with an inert anode. In industry, oxygen is obtained by separating liquid air in distillation columns.

CHEMICAL THERMODYNAMICS

Thermodynamics is the science of the interconversions of various forms of energy and the laws of these transformations. Thermodynamics is based only on experimentally discovered objective laws expressed in the two basic principles of thermodynamics.

Thermodynamics studies:

1. Transitions of energy from one form to another, from one part of the system to another;

2. Energy effects accompanying various physical and chemical processes and their dependence on the conditions of these processes;

3. Possibility, direction and limits of spontaneous occurrence of processes under the conditions under consideration.

It should be noted that classical thermodynamics has the following limitations:

1. Thermodynamics does not consider the internal structure of bodies and the mechanism of processes occurring in them;

2. Classical thermodynamics studies only macroscopic systems;

3. In thermodynamics there is no concept of “time”.

BASIC CONCEPTS OF THERMODYNAMICS

A thermodynamic system is a body or group of bodies interacting, mentally or actually isolated from the environment.

A homogeneous system is a system within which there are no surfaces separating parts of the system (phases) that differ in properties.

A heterogeneous system is a system within which there are surfaces that separate parts of the system that differ in properties.

A phase is a collection of homogeneous parts of a heterogeneous system, identical in physical and chemical properties, separated from other parts of the system by visible interfaces.

An isolated system is a system that does not exchange either matter or energy with the environment.

A closed system is a system that exchanges energy with the environment, but does not exchange matter.

An open system is a system that exchanges both matter and energy with the environment.

Components of a system are individual substances that, when taken in the smallest quantity, are sufficient to describe (form) all phases of the system. the release of components is determined by the specific content of the system and depends on the chemical reactions that occur within the system and during its interaction with the external environment. In complex mineral systems, oxides or elements usually act as components.

Parameters are quantities that can be used to describe the state of a system. Fundamental parameters of systems: temperature (T), entropy (S), pressure (p), volume (V), masses of components (m a ...m k) and their chemical potentials (μ a ...μ k).

Extensive parameters are those that have the property of additivity (composition), i.e. extensive parameters depend on the mass or number of particles of the system. Extensive parameters include volume, entropy and masses of components. Extensive parameters are sometimes called capacity parameters. Intensive parameters, or intensity parameters, are those that do not depend on the mass or number of particles of the system. These include temperature, pressure and chemical potentials of the components.

There is a remarkable property of thermodynamic parameters, which can be called the property of symmetry and conjugacy. The property of symmetry is that any thermodynamic process in a system is characterized by a pair of parameters, one of which is intensive, the other extensive.

The first law of thermodynamics is the law of conservation of energy, one of the universal laws of nature: Energy is indestructible and uncreated; it can only pass from one form to another in equivalent proportions.

The first law of thermodynamics is the postulate

The total energy of an isolated system is constant;

A perpetual motion machine of the first kind (an engine that does work without expending energy) is impossible.

The first law of thermodynamics establishes the relationship between heat Q, work A and the change in internal energy of the system ΔU: Equation 1 is a mathematical representation of the 1st law of thermodynamics for a finite, equation 2 for an infinitesimal change in the state of the system.

Internal energy is a function of state; this means that the change in internal energy ΔU does not depend on the path of transition of the system from state 1 to state 2 and is equal to the difference between the values of internal energy U 2 and U 1 in these states:

Isochoric process (V = const; ΔV = 0). Absorption or release of heat is associated only with the release of E

An isothermal process (T = const). This is a process of quasi-static expansion or compression of a substance in contact with a thermal reservoir.

Isobaric process (P = const).

Adiabatic process (Q = 0). This is the process of quasi-static expansion or compression of a gas in a vessel with heat-tight shades. A=- U

INTERNAL ENERGY thermodynamic. function of the state of the system, its energy, determined internally. condition. Internal energy is basically added up. from kinetic energy of movement of particles (atoms, molecules, ions, electrons) and energy of interaction. between them (intra- and intermolecular).

During an isothermal process, the internal energy of an ideal gas does not change. The entire amount of heat transferred to the gas is used to perform work: Q = A

Change in internal energy during an isobaric process: ΔU=3/2 ·v·R·ΔT.

change in internal energy at adiabatic: Q=m·C p D·T/m.

Enthalpy is a quantity proportional to the quantity of a substance and is measured in [KJ/mol] N<0-экзотермический, Н>0 endothermic.

When interacting with gases H2O is formed, which can be in different states.

Standard enthalpy state T=298K, P=101.325kPa