Why does sulfur not dissolve in water? Sulfur compounds. Preparation of sulfuric acid
Description and properties of sulfur
Sulfur is a substance that is in group 16, under the third period and has an atomic number of 16. It can be found both in native and bound form. Sulfur is designated by the letter S. Known sulfur formula– (Ne)3s 2 3p 4 . Sulfur as an element is included in many proteins.
The photo shows sulfur crystals
If speak about atomic structure of the element sulfur, then in its outer orbit there are electrons whose valence number reaches six.
This explains the element's property of being maximally hexavalent in most combinations. There are four isotopes in the structure of a natural chemical element, and these are 32S, 33S, 34S and 36S. Speaking about the outer electron shell, the atom has a 3s2 3p4 scheme. The radius of the atom is 0.104 nanometers.
Properties of sulfur are primarily divided into physical types. This includes the fact that the element has a solid crystalline composition. Two allotropic modifications are the main state in which this sulfur element is stable.
The first modification is rhombic, lemon-yellow in color. Its stability is lower than 95.6 °C. The second is monoclinic, having a honey-yellow color. Its resistance ranges from 95.6 °C and 119.3 °C.
The photo shows the mineral sulfur
During smelting, the chemical element becomes a moving liquid that is yellow in color. It turns brown, reaching temperatures of more than 160 °C. And at 190 °C sulfur color turns into dark brown. After reaching 190 °C, a decrease in the viscosity of the substance is observed, which nevertheless becomes liquid after heating to 300 °C.
Other properties of sulfur:
Practically does not conduct heat or electricity.
Does not dissolve when immersed in water.
It is soluble in ammonia, which has an anhydrous structure.
It is also soluble in carbon disulfide and other organic solvents.
TO characteristics of the element sulfur it is important to add its chemical features. She is active in this regard. If sulfur is heated, it can simply combine with almost any chemical element.
The photo shows a sample of sulfur mined in Uzbekistan
With the exception of inert gases. Upon contact with metals, chemicals. the element forms sulfides. Room temperature allows the element to react with. Increased temperature increases the activity of sulfur.
Let's consider how sulfur behaves with individual substances:
With metals it is an oxidizing agent. Forms sulfides.
Active interaction occurs with hydrogen at high temperatures – up to 200 °C.
With oxygen. Oxides form at temperatures up to 280 °C.
With phosphorus, carbon – it is an oxidizing agent. Only if there is no air during the reaction.
With fluorine it acts as a reducing agent.
With substances that have a complex structure - also as a reducing agent.
Sulfur deposits and production
The main source for obtaining sulfur is its deposits. In total, there are 1.4 billion tons of reserves of this substance worldwide. It is mined both by open and underground mining and by smelting from underground.
The photo shows sulfur mining in the Kawa Ijen volcano
If the latter case applies, then water is used, which is overheated and melts the sulfur with it. In low-grade ores, the element is contained in approximately 12%. Rich – 25% and more.
Common types of deposits:
Stratiform – up to 60%.
Salt dome - up to 35%.
Volcanogenic – up to 5%.
The first type is associated with strata called sulfate-carbonate. At the same time, ore bodies that have a thickness of up to several tens of meters and a size of up to hundreds of meters are located in sulfate rocks.
Also, these strata deposits can be found among rocks of sulfate and carbonate origin. The second type is characterized by gray deposits, which are confined to salt domes.
The latter type is associated with volcanoes that have a young and modern structure. In this case, the ore element has a sheet-like, lens-shaped shape. It may contain sulfur in the amount of 40%. This type of deposit is common in the Pacific volcanic belt.
Sulfur deposit in Eurasia is located in Turkmenistan, the Volga region and other places. Sulfur rocks are found near the left banks of the Volga, which stretch from Samara. The width of the rock strip reaches several kilometers. Moreover, they can be found all the way to Kazan.
The photo shows sulfur in rock
In Texas and Louisiana, huge amounts of sulfur are found in the roofs of salt domes. Particularly beautiful Italians of this element are found in Romagna and Sicily. And on the island of Vulcano they find monoclinic sulfur. The element, which was oxidized by pyrite, was found in the Urals in the Chelyabinsk region.
For mining sulfur chemical element use different methods. It all depends on the conditions of its occurrence. At the same time, of course, special attention is paid to safety.
Since hydrogen sulfide accumulates along with sulfur ore, it is necessary to take a particularly serious approach to any mining method, because this gas is poisonous to humans. Sulfur also tends to ignite.
Most often they use the open method. So, with the help of excavators, significant parts of the rocks are removed. Then the ore part is crushed using explosions. The lumps are sent to the factory for enrichment. Then - to the sulfur smelting plant, where sulfur is obtained from concentrate.
The photo shows sulfur in the port, brought by sea
In the case of deep occurrence of sulfur in many volumes, the Frasch method is used. The sulfur melts while still underground. Then, like oil, it is pumped out through a broken well. This approach is based on the fact that the element melts easily and has a low density.
A separation method using centrifuges is also known. Only this method has a drawback: sulfur is obtained with impurities. And then it is necessary to carry out additional cleaning.
In some cases, the borehole method is used. Other possibilities for mining the sulfur element:
Steam-water.
Filtration.
Thermal.
Centrifugal.
Extraction.
Application of sulfur
Most of the mined sulfur is used to make sulfuric acid. And the role of this substance is very huge in chemical production. It is noteworthy that to obtain 1 ton of sulfuric substance, 300 kg of sulfur is needed.
Sparklers, which glow brightly and have many dyes, are also made using sulfur. The paper industry is another area where a significant portion of the extracted substance goes.
Pictured is sulfur ointment
More often application of sulfur finds when meeting production needs. Here are some of them:
Use in chemical production.
For the production of sulfites, sulfates.
Production of substances for fertilizing plants.
To obtain non-ferrous types of metals.
To give steel additional properties.
For making matches, materials for explosions and pyrotechnics.
Paints and fibers from artificial materials are produced using this element.
For bleaching fabrics.
In some cases sulfur element included in ointments that treat skin diseases.
Sulfur price
According to the latest news, the need for sulfur is actively growing. The cost of a Russian product is 130 dollars. For the Canadian version – $145. But in the Middle East, prices increased to $8, resulting in a cost of $149.
The photo shows a large specimen of the mineral sulfur
In pharmacies you can find ground sulfur powder at a price of 10 to 30 rubles. In addition, it is possible to buy it in bulk. Some organizations offer to purchase granular technical equipment at a low price. gas sulfur.
The content of the article
SULFUR, S (sulfur), a non-metallic chemical element, a member of the chalcogen family (O, S, Se, Te and Po) – Group VI of the periodic table of elements. Sulfur, like many of its uses, have been known since ancient times. A. Lavoisier argued that sulfur is an element. Sulfur is vital for the growth of plants and animals; it is part of living organisms and their decomposition products; there is a lot of it, for example, in eggs, cabbage, horseradish, garlic, mustard, onions, hair, wool, etc. It is also present in coals and oil.
Application.
About half of annual sulfur consumption goes into the production of industrial chemicals such as sulfuric acid, sulfur dioxide and carbon disulfide (carbon disulfide). In addition, sulfur is widely used in the production of insecticides, matches, fertilizers, explosives, paper, polymers, paints and dyes, and in the vulcanization of rubber. The leading place in sulfur production is occupied by the USA, CIS countries and Canada.
Prevalence in nature.
Sulfur occurs in a free state (native sulfur). In addition, there are huge reserves of sulfur in the form of sulfide ores, primarily the ores of lead (lead luster), zinc (zincblende), copper (copper luster) and iron (pyrite). When metals are extracted from these ores, they are freed from sulfur, usually by roasting in the presence of oxygen, which produces sulfur dioxide (IV), which is often released into the atmosphere without use. In addition to sulfide ores, quite a lot of sulfur is found in the form of sulfates, for example, calcium sulfate (gypsum), barium sulfate (barite). Sea water and many mineral waters contain water-soluble magnesium and sodium sulfates. Hydrogen sulfide (hydrogen sulfide) is found in some mineral waters. In industry, sulfur can be obtained as a by-product of processes in smelters, coke ovens, oil refining, from flue gases or natural gases. Sulfur is extracted from natural underground deposits by melting it with superheated water and delivering it to the surface using compressed air and pumps. In the flash process for extracting sulfur from sulfur deposits using a concentric pipe installation, patented by G. Frasch in 1891, sulfur is obtained with a purity of up to 99.5%.
Properties.
Sulfur has the form of a yellow powder or brittle crystalline mass, odorless and tasteless and insoluble in water. Sulfur is characterized by several allotropic modifications. The most famous are the following: crystalline sulfur - orthorhombic (native sulfur, a-S) and monoclinic (prismatic sulfur, b-S); amorphous - colloidal (sulfur milk) and plastic; intermediate amorphous-crystalline - sublimated (sulfur color).
Crystalline sulfur.
Crystalline sulfur has two modifications; one of them, orthorhombic, is obtained from a solution of sulfur in carbon disulfide (CS 2) by evaporating the solvent at room temperature. In this case, diamond-shaped translucent crystals of light yellow color are formed, easily soluble in CS 2. This modification is stable up to 96°C; at higher temperatures the monoclinic form is stable. With the natural cooling of molten sulfur in cylindrical crucibles, large crystals of the orthorhombic modification with a distorted shape (octahedra with corners or faces partially “cut off”) grow. This material is called lump sulfur in industry. The monoclinic modification of sulfur is long transparent dark yellow needle-shaped crystals, also soluble in CS 2. When monoclinic sulfur is cooled below 96° C, a more stable yellow orthorhombic sulfur is formed.
Non-crystalline sulfur.
In addition to these crystalline and amorphous forms, there is an intermediate form known as sulfur color or sublimated sulfur, which is produced by the condensation of sulfur vapor without passing through the liquid phase. It consists of tiny grains with a crystallization center and an amorphous surface. This form dissolves slowly and incompletely in CS 2 . After treatment with ammonia to remove impurities such as arsenic, the product is known medically as washed sulfur, which is used in a similar way to colloidal sulfur.
Liquid state.
Sulfur molecules consist of a closed chain of eight atoms (S 8). Liquid sulfur has an unusual property: with increasing temperature, its viscosity increases. Below 160° C, sulfur is a typical yellowish liquid, its composition corresponds to the formula S 8 and is designated l-S. As the temperature rises, the ring molecules of S8 begin to break and join together, forming long chains ( m-S), the color of liquid sulfur becomes dark red, viscosity increases, reaching a maximum at 200–250 ° C. With a further increase in temperature, liquid sulfur becomes lighter, long chains break, forming short ones, with less ability to intertwine, which leads to lower viscosity.
Gas.
Sulfur boils at 444.6° C, forming orange-yellow vapors consisting mainly of S 8 molecules. With increasing temperature, the color of the vapor turns into dark red, then into fawn, and at 650° C into straw yellow. With further heating, the S 8 molecules dissociate, forming the equilibrium forms S 6, S 4 and S 2 at different temperatures. And finally, at >1000° C the vapors consist practically of S 2 molecules, and at 2000° C they consist of monatomic molecules.
Chemical properties.
Sulfur is a typical non-metal. It has six electrons on its outer electron shell, and it more easily attaches electrons to other elements than gives up its own. It reacts with many metals, releasing heat (for example, when combined with copper, iron, zinc). It also combines with almost all non-metals, although not as vigorously.
Connections.
Sulfur dioxide
is formed when sulfur is burned in air, in particular during the roasting of sulfide metal ores. Sulfur dioxide is a colorless gas with a suffocating odor. It is an anhydride of sulfurous acid and readily dissolves in water to form sulfurous acid. The dioxide easily liquefies (boiling point -10°C) and is stored in steel cylinders. Dioxide is used in the production of sulfuric acid, in refrigeration units, for bleaching textiles, wood pulp, straw, beet sugar, for preserving fruits and vegetables, for disinfection, in brewing and food production.
Sulfurous acid
H 2 SO 3 exists only in dilute solutions (less than 6%). It is a weak acid that forms medium and acid salts (sulfites and hydrosulfites). Sulfurous acid is a good reducing agent and reacts with oxygen to form sulfuric acid. Sulfurous acid has several applications, including bleaching silk, wool, paper, wood pulp and similar substances. It is used as an antiseptic and preservative, especially to prevent the fermentation of wine in barrels, to prevent the fermentation of grains during the extraction of starch. Acid is also used to preserve food. The most important of its salts is calcium hydrosulfite Ca(HSO 3) 2, used in the processing of wood chips into cellulose.
Sulfur trioxide
SO 3 (sulfuric anhydride), which forms sulfuric acid with water, is either a colorless liquid or a white crystalline substance (crystallizes at 16.8 ° C; bp 44.7 ° C). It is formed by the oxidation of sulfur dioxide with oxygen in the presence of an appropriate catalyst (platinum, vanadium pentoxide). Sulfur trioxide smokes strongly in moist air and dissolves in water, forming sulfuric acid and generating a lot of heat. It is used in the production of sulfuric acid and the production of synthetic organic substances.
Sulfuric acid
H2SO4. Anhydrous H 2 SO 4 is a colorless oily liquid that dissolves SO 3 to form oleum. Miscible with water in any ratio. When dissolved in water, hydrates are formed with the release of a very large amount of heat; therefore, in order to avoid splashing of the acid, it is usually necessary to carefully add the acid to the water when dissolving, and not vice versa. Concentrated acid absorbs water vapor well and is therefore used to dry gases. For the same reason, it leads to charring of organic substances, especially carbohydrates (starch, sugar, etc.). If it comes into contact with the skin, it causes severe burns; the vapors corrode the mucous membranes of the respiratory tract and eyes. Sulfuric acid is a strong oxidizing agent. Conc. H 2 SO 4 oxidizes HI, HBr to I 2 and Br 2, respectively, coal to CO 2, sulfur to SO 2, metals to sulfates. Dilute acid also oxidizes metals in the voltage series up to hydrogen. H 2 SO 4 is a strong dibasic acid that forms medium and acid salts - sulfates and hydrosulfates; Most of its salts are soluble in water, with the exception of barium, strontium and lead sulfates; calcium sulfate is slightly soluble.
Sulfuric acid is one of the most important products of the chemical industry (producing alkalis, acids, salts, mineral fertilizers, chlorine). It is obtained mainly by contact or tower method according to the following principle:
Most of the resulting acid is used in the production of mineral fertilizers (superphosphate, ammonium sulfate). Sulfuric acid serves as a starting material for the production of salts and other acids, for the synthesis of organic substances, artificial fibers, for the purification of kerosene, petroleum oils, benzene, toluene, in the manufacture of paints, etching of ferrous metals, in the hydrometallurgy of uranium and some non-ferrous metals, for the production of detergents and medicines, as an electrolyte in lead batteries and as a desiccant.
Thiosulfuric acid
H 2 S 2 O 3 is structurally similar to sulfuric acid except for the replacement of one oxygen with a sulfur atom. The most important derivative of the acid is sodium thiosulfate Na 2 S 2 O 3 - colorless crystals formed by boiling sodium sulfite Na 2 SO 3 with a sulfur color. Sodium thiosulfate (or hyposulfite) is used in photography as a fixative.
Sulfonal
(CH 3) 2 C(SO 2 C 2 H 5) 2 is a white crystalline substance, odorless, slightly soluble in water, is a narcotic and is used as a sedative and hypnotic.
Hydrogen sulfide
H 2 S (hydrogen sulfide) is a colorless gas with a pungent, unpleasant odor of rotten eggs. It is slightly heavier than air (density 1.189 g/dm3), easily liquefies into a colorless liquid and is highly soluble in water. The solution in water is a weak acid with a pH of ~ 4. Liquid hydrogen sulfide is used as a solvent. The solution and gas are widely used in qualitative analysis for the separation and determination of many metals. Inhalation of small amounts of hydrogen sulfide causes headaches and nausea, large amounts or continuous inhalation of hydrogen sulfide causes paralysis of the nervous system, heart and lungs. Paralysis occurs unexpectedly, as a result of disruption of the vital functions of the body.
Sulfur monochloride
S 2 Cl 2 is a fuming, amber-colored oily liquid with a pungent odor, tearing and making breathing difficult. It smokes in moist air and decomposes with water, but is soluble in carbon disulfide. Sulfur monochloride is a good solvent for sulfur, iodine, metal halides and organic compounds. The monochloride is used for the vulcanization of rubber, in the production of printing ink and insecticides. When reacted with ethylene, a volatile liquid is formed known as mustard gas (ClC 2 H 4) 2 S, a toxic compound used as a chemical irritant.
Carbon disulfide
CS 2 (carbon disulfide) is a pale yellow liquid, toxic and flammable. CS 2 is produced by synthesis from elements in an electric furnace. The substance is insoluble in water, has a high refractive index, high vapor pressure, and low boiling point (46° C). Carbon disulfide - an effective solvent for fats, oils, rubber and rubber - is widely used for the extraction of oils, in the production of artificial silk, varnishes, rubber adhesives and matches, the destruction of barn weevils and clothing moths, and for soil disinfection.
Sulfur(S) is a chemical element of group 16 of the periodic system of elements with atomic number 16, the simple substance of which sulfur - non-metal, yellow crystalline substance. It is found in nature in the native state and in the form of heavy metal sulfides (pyrite and others). Sulfur is used primarily in the chemical industry for the production of sulfuric acid, synthetic fiber, sulfur dyes, black powder, in the rubber industry, as well as in agriculture, pharmaceuticals, etc.
Due to its ability to form disulfide bonds, sulfur plays an important role in the composition of proteins.
Story
The elemental nature of sulfur was established by Antoine Lavoisier in his combustion experiments.
general characteristics
Sulfur has an atomic mass of 32.06. In nature, there are 4 stable isotopes with mass numbers 32-34 and 36. Sulfur belongs to the chalcogens, according to the new classification in the sixteenth, and according to the old one to the VI group of elements of the periodic table. Sulfur is a non-metal.
Several allotropic forms of sulfur are known. Under normal conditions, rhombic sulfur is stable - pale yellow in color, with a density of 2070 kg / m3, t melt = 112.8 ° C, t boil = 444.6 o C. In all liquid and solid states, sulfur is diamagnetic. The thermodynamic and other properties of sulfur change sharply at 160 °C, which is associated with a change in the molecular structure of liquid sulfur. The viscosity of sulfur increases greatly with increasing temperature (from 0.0065 Pas at 155 °C to 93.3 Pas at 187 °C), and then falls (to 0.083 Pas at 444.6 °C).
Sulfur reacts with almost all metals.
Distribution in nature
Sulfur is a fairly common element, accounting for about 0.1% of the mass of the earth's crust. The average sulfur content in the earth's crust is 4.710 -2 wt.%, while the main amount of natural sulfur is concentrated in sedimentary rocks (0.3 wt.%). In other rocks, the average sulfur content is as follows: dunites, peridotites, pyroxenites - 0.01%; basalts, habronorites, diabase - 0.03%; diorites, andesites - 0.02%.
In nature, sulfur is found both in a free state - the so-called native sulfur, but much more often it is found in a bound form, that is, in the form of various compounds. The most important of them are iron pyrite, or pyrite FeS 2, zinc blende ZnS, lead luster PbS, copper luster Cu 2 S, gypsum CaSO 4 2H 2 O, mirabilite Na 2 SO 4 10H 2 O, etc.
Sulfur is found in coal and oil, as well as in all plant and animal organisms, since it is part of proteins.
The sulfur content of oil and natural gas is estimated at 210 9 tons, that is, more than natural sulfur reserves. Sulfur in oil is present in different forms, from elemental sulfur and hydrogen sulfide to organic sulfur, which includes more than 120 compounds. The main sulfur-containing substances of hydrocarbon raw materials are hydrogen sulfide, mercaptans and other organosulfur compounds. The raw material base for the production of sulfur is, as a rule, gases containing at least 0.1% hydrogen sulfide.
Of course, native sulfur is found in a continuous mass, filling cracks and cavities in rocks, or in the form of sinter, spherical and nest-like aggregates, stalactites, stalagmites, deposits, efflorescences, and earthy powdery accumulations. It often forms crystals, which are often grouped into intergrowths, druses, and brushes.
Physical properties
Sulfur is a yellow crystalline substance. It is very fragile and easily ground into the smallest powder. Density 2070 kg/m3. t melt = 112.8 ° C, t boil = 444.6 o C. In all liquid and solid states, sulfur is diamagnetic.
It is found in three allotropic forms: two crystalline (orthorhombic and monoclinic, according to the method of joining atoms in the crystal) and amorphous.
- α-S (orthorhombic) crystal modification, t melt = 112.8 ° C, stable at 95.6 ° C, lemon yellow;
- β-S crystal modification, t melt = 119 ° C, stable at 95.6-119 ° C, honey-yellow. Up to 160 ° C, the molecules are 8-atomic, in pairs - 2-atomic (paramagnetic sulfur), 4, 6, and 8-atomic.
- Above 160 °C, helical chains of μ-S plastic sulfur are formed.
Sulfur almost does not conduct electric current or heat. When cooled very quickly, sulfur vapor turns into a solid state in the form of a very fine powder (sulfur-colored), bypassing the liquid state. Sulfur is insoluble in water and is not wetted by water, but it dissolves well in benzene C 6 H 6 and especially in carbon disulfide CS 2.
Chemical properties
Having six electrons in the outer layer: (+ 16), 2,8,6 - sulfur atoms exhibit the properties of an oxidizing agent and, by adding two electrons from the atoms of other elements, which they lack in a completely filled outer shell, they turn into negative divalent ions: S 0 + 2e = S 2. But Sulfur is a less active oxidizing agent than oxygen, since its valence electrons are distant from the nucleus of the atom and are weaker bound to it than the valence electrons of oxygen atoms. Unlike oxygen, Sulfur can also exhibit the properties of a reducing agent: S 0 - 6e = S 6+ or S 0 - 4e = S 4+. The reducing properties of sulfur manifest themselves when interacting with a stronger oxidizing agent, that is, with substances whose atoms have a greater affinity for electrons.
Sulfur can react directly with almost all metals (with the exception of noble metals), but mainly when heated. So, if a mixture of sulfur and iron powders is heated at least in one place so that a reaction begins, then the entire mixture will heat up by itself (due to the heat of reaction) and turn into a black, brittle substance - iron monosulfide:
Fe + S = FeS
When ignited, a mixture of sulfur and zinc powders reacts very violently, with a flash. As a result of the reaction, zinc sulfide is formed:
Zn + S = ZnS
Sulfur reacts with mercury even at ordinary temperatures. Thus, when mercury is ground with sulfur powder, a black substance appears - mercury sulfide:
Hg + S = HgS
At high temperatures, sulfur also reacts with hydrogen to form hydrogen sulfide:
H 2 + S = H 2 S.
When interacting with metals and hydrogen, sulfur plays the role of an oxidizing agent, and itself is reduced to S 2 ions. Therefore, in all sulfides, sulfur is negatively divalent. Sulfur also reacts relatively easily with oxygen. Thus, ignited sulfur burns in air to form sulfur dioxide SO 2 (sulfite anhydride) and a very small amount of sulfur trioxide SO 3 (sulfate anhydride).
- S + O 2 = SO 2
- 2S + 3O 2 = 2SO 3
In this case, oxygen is the oxidizing agent, and sulfur is the reducing agent. In the first reaction, the sulfur atom loses four, and in the second, six valence electrons, as a result of which Sulfur in SO 2 is positively tetravalent, and in SO 3 it is positively hexavalent.
Receipt
Sulfur is obtained from native ores, as well as as a by-product during the processing of polymetallic ores, from sulfates during their complex processing, from natural gases and fossil fuels during their purification. The proportion of sulfur obtained from hydrogen sulfide increases. To separate sulfur from foreign impurities, it is smelted in autoclaves. Autoclaves are iron cylinders into which ore is loaded and heated with superheated steam to 150 ° C under a pressure of 6 atm. The molten sulfur flows down, and the waste rock remains. Smelted from ore, sulfur still contains a certain amount of impurities.
Completely pure sulfur is obtained by distillation in special furnaces connected to large chambers. Sulfur vapor in a cold chamber immediately turns into a solid state and settles on the walls in the form of a very fine light yellow powder. When the chamber heats up to 120 ° C, the sulfur vapor turns into liquid. Molten sulfur is poured into wooden cylindrical molds, where it hardens. Such sulfur is called Cherenkova.
Application
Sulfur is widely used in various sectors of the national economy, mainly in the chemical industry for the production of sulfuric acid H 2 SO 4 (almost half of the sulfur produced in the world), carbon disulfide CS 2, some dyes, and other chemical products. Significant amounts of sulfur are consumed by the rubber industry to vulcanize rubber, that is, to convert rubber into rubber.
Sulfur is used in the chemical industry in the production of phosphoric, hydrochloric and other acids, in the rubber industry, in the production of dyes, black powder, and the like. Native sulfur is used in agriculture (insecticides, microfertilizers, as a disinfectant in livestock farming).
Technical sulfur used for the production of sulfuric acid must contain at least 95% sulfur, arsenic and selenium should not be present at all, and the content of organic substances should not exceed 1%. Man-made fiber (viscose) production in the chemical industry is another consumer of sulfur. In agriculture, sulfur is used as a means of pest control, partly as a fertilizer, and for disinfection in the treatment of animals. In paper production, sulfur in the form of SO2 is used when processing wood pulp (bisulfate method). Sulfur is used in rubber vulcanization, glass, and leather industries. Minor amounts of high purity sulfur are used in the chemical and pharmaceutical industries. Sulfur is also used to produce ultramarine. The textile, food, starch and molasses industries use sulfur or its compounds for bleaching and clarification, in fruit canning, and in refrigeration.
Sulfur is also used in match production, in pyrotechnics, in the production of black powder, and the like. In medicine, sulfur is used to make sulfur ointment for the treatment of skin diseases. In agriculture, sulfur is used to control pests of cotton and grapevines.
Impact on humans
Sulfur dust irritates the respiratory system and mucous membranes. MPC – 2 mg/m3.
Sulfur– element of the 3rd period and VIA group of the Periodic System, serial number 16, refers to chalcogens. The electronic formula of the atom is [ 10 Ne]3s 2 3p 4, the characteristic oxidation states are 0, -II, +IV and +VI, the S VI state is considered stable.
Scale of sulfur oxidation states:
The electronegativity of sulfur is 2.60 and is characterized by non-metallic properties. In hydrogen and oxygen compounds it is found in various anions and forms oxygen-containing acids and their salts, binary compounds.
In nature - fifteenth element by chemical abundance (seventh among non-metals). It is found in free (native) and bound form. A vital element for higher organisms.
Sulfur S. Simple substance. Yellow crystalline (α‑rhombic and β‑monoclinic,
at 95.5 °C) or amorphous (plastic). At the nodes of the crystal lattice there are S 8 molecules (non-planar rings of the “crown” type), amorphous sulfur consists of S n chains. A low-melting substance, the viscosity of the liquid passes through a maximum at 200 °C (breakdown of S 8 molecules, interweaving of S n chains). The pair contains molecules S 8, S 6, S 4, S 2. At 1500 °C, monoatomic sulfur appears (in chemical equations, for simplicity, any sulfur is depicted as S).
Sulfur is insoluble in water and under normal conditions does not react with it; it is highly soluble in carbon disulfide CS 2.
Sulfur, especially powdered sulfur, is highly active when heated. Reacts as an oxidizing agent with metals and non-metals:
but as reducing agent– with fluorine, oxygen and acids (boiling):
Sulfur undergoes dismutation in alkali solutions:
3S 0 + 6KOH (conc.) = 2K 2 S ‑II + K 2 S IV O 3 + 3H 2 O
At high temperatures (400 °C), sulfur displaces iodine from hydrogen iodide:
S + 2HI (g) = I 2 + H 2 S,
but in solution the reaction goes in the opposite direction:
I 2 + H 2 S (p) = 2 HI + S↓
Receipt: V industry smelted from natural deposits of native sulfur (using water vapor), released during desulfurization of coal gasification products.
Sulfur is used for the synthesis of carbon disulfide, sulfuric acid, sulfur (vat) dyes, in the vulcanization of rubber, as a means of protecting plants from powdery mildew, and for the treatment of skin diseases.
Hydrogen sulfide H 2 S. Anoxic acid. A colorless gas with a suffocating odor, heavier than air. The molecule has the structure of a doubly incomplete tetrahedron [::S(H) 2 ]
(sp 3 -hybridization, valet angle H – S–H is far from tetrahedral). Unstable when heated above 400 °C. Slightly soluble in water (2.6 l/1 l H 2 O at 20 °C), saturated decimolar solution (0.1 M, “hydrogen sulfide water”). A very weak acid in solution, practically does not dissociate in the second stage to S 2‑ ions (the maximum concentration of S 2‑ is 1 10 ‑ 13 mol/l). When exposed to air, the solution becomes cloudy (the inhibitor is sucrose). Neutralized by alkalis, but not completely by ammonia hydrate. Strong reducing agent. Enters into ion exchange reactions. A sulfiding agent precipitates differently colored sulfides with very low solubility from solution.
Qualitative reactions– precipitation of sulfides, as well as incomplete combustion of H 2 S with the formation of a yellow sulfur deposit on a cold object brought into the flame (porcelain spatula). A by-product of oil, natural and coke oven gas refining.
It is used in the production of sulfur, inorganic and organic sulfur-containing compounds as an analytical reagent. Extremely poisonous. Equations of the most important reactions:
Receipt: V industry– direct synthesis:
H 2 + S = H2S(150–200 °C)
or by heating sulfur with paraffin;
V laboratories– displacement from sulfides with strong acids
FeS + 2НCl (conc.) = FeCl 2 + H2S
or complete hydrolysis of binary compounds:
Al 2 S 3 + 6H 2 O = 2Al(OH) 3 ↓ + 3 H2S
Sodium sulfide Na 2 S. Oxygen-free salt. White, very hygroscopic. Melts without decomposition, thermally stable. It is highly soluble in water, hydrolyzes at the anion, and creates a highly alkaline environment in solution. When exposed to air, the solution becomes cloudy (colloidal sulfur) and turns yellow (polysulfide color). Typical reducer. Adds sulfur. Enters into ion exchange reactions.
Qualitative reactions on the S 2‑ ion – precipitation of differently colored metal sulfides, of which MnS, FeS, ZnS decompose into HCl (diluted).
It is used in the production of sulfur dyes and cellulose, for removing hair from hides when tanning leather, as a reagent in analytical chemistry.
Equations of the most important reactions:
Na 2 S + 2НCl (diluted) = 2NaCl + H 2 S
Na 2 S + 3H 2 SO 4 (conc.) = SO 2 + S↓ + 2H 2 O + 2NaHSO 4 (up to 50 °C)
Na 2 S + 4HNO 3 (conc.) = 2NO + S↓ + 2H 2 O + 2NaNO 3 (60 °C)
Na 2 S + H 2 S (saturated) = 2NaHS
Na 2 S (t) + 2O 2 = Na 2 SO 4 (above 400 °C)
Na 2 S + 4H 2 O 2 (conc.) = Na 2 SO 4 + 4H 2 O
S 2‑ + M 2+ = MnS (tel.)↓; FeS (black)↓; ZnS (white)↓
S 2‑ + 2Ag + = Ag 2 S (black)↓
S 2‑ + M 2+ = СdS (yellow)↓; PbS, CuS, HgS (black)↓
3S 2‑ + 2Bi 3+ = Bi 2 S 3 (cor. – black)↓
3S 2‑ + 6H 2 O + 2M 3+ = 3H 2 S + 2M(OH) 3 ↓ (M = Al, Cr)
Receipt V industry– calcination of the mineral mirabilite Na 2 SO 4 10H 2 O in the presence of reducing agents:
Na 2 SO 4 + 4H 2 = Na 2 S + 4H 2 O (500 °C, cat. Fe 2 O 3)
Na 2 SO 4 + 4С (coke) = Na 2 S + 4СО (800–1000 °C)
Na 2 SO 4 + 4СО = Na 2 S + 4СО 2 (600–700 °C)
Aluminum sulfide Al 2 S 3. Oxygen-free salt. White, the Al–S bond is predominantly covalent. Melts without decomposition under excess pressure N 2, easily sublimes. Oxidizes in air when heated. It is completely hydrolyzed by water and does not precipitate from solution. Decomposes with strong acids. Used as a solid source of pure hydrogen sulfide. Equations of the most important reactions:
Al 2 S 3 + 6H 2 O = 2Al(OH) 3 ↓ + 3H 2 S (pure)
Al 2 S 3 + 6HCl (diluted) = 2AlCl 3 + 3H 2 S
Al 2 S 3 + 24HNO 3 (conc.) = Al 2 (SO 4) 3 + 24NO 2 + 12H 2 O (100 °C)
2Al 2 S 3 + 9O 2 (air) = 2Al 2 O 3 + 6SO 2 (700–800 °C)
Receipt: interaction of aluminum with molten sulfur in the absence of oxygen and moisture:
2Al + 3S = AL 2 S 3(150–200 °C)
Iron (II) sulfide FeS. Oxygen-free salt. Black-gray with a green tint, refractory, decomposes when heated in a vacuum. When wet, it is sensitive to air oxygen. Insoluble in water. Does not precipitate when solutions of iron(II) salts are saturated with hydrogen sulfide. Decomposes with acids. It is used as a raw material in the production of cast iron, a solid source of hydrogen sulfide.
The iron(III) compound Fe 2 S 3 is not known (not obtained).
Equations of the most important reactions:
Receipt:
Fe + S = FeS(600 °C)
Fe 2 O 3 + H 2 + 2H 2 S = 9 FeS+ 3H 2 O (700‑1000 °C)
FeCl 2 + 2NH 4 HS (g) = FeS↓ + 2NH 4 Cl + H 2 S
Iron disulfide FeS 2. Binary connection. It has the ionic structure Fe 2+ (–S – S–) 2‑ . Dark yellow, thermally stable, decomposes when heated. Insoluble in water, does not react with dilute acids and alkalis. Decomposes by oxidizing acids and is fired in air. It is used as a raw material in the production of cast iron, sulfur and sulfuric acid, and a catalyst in organic synthesis. Ore minerals found in nature pyrite And Marcasite.
Equations of the most important reactions:
FeS 2 = FeS + S (above 1170 °C, vacuum)
2FeS 2 + 14H 2 SO 4 (conc., horizontal) = Fe 2 (SO 4) 3 + 15SO 2 + 14H 2 O
FeS 2 + 18HNO 3 (conc.) = Fe(NO 3) 3 + 2H 2 SO 4 + 15NO 2 + 7H 2 O
4FeS 2 + 11O 2 (air) = 8SO 2 + 2Fe 2 O 3 (800 °C, roasting)
Ammonium hydrosulfide NH 4 HS. An oxygen-free acidic salt. White, melts under excess pressure. Very volatile, thermally unstable. It oxidizes in air. It is highly soluble in water, hydrolyzes into the cation and anion (predominates), creates an alkaline environment. The solution turns yellow in air. Decomposes with acids and adds sulfur in a saturated solution. It is not neutralized by alkalis, the average salt (NH 4) 2 S does not exist in solution (for the conditions for obtaining the average salt, see the section “H 2 S”). It is used as a component of photographic developers, as an analytical reagent (sulfide precipitator).
Equations of the most important reactions:
NH 4 HS = NH 3 + H 2 S (above 20 °C)
NH 4 HS + HCl (diluted) = NH 4 Cl + H 2 S
NH 4 HS + 3HNO 3 (conc.) = S↓ + 2NO 2 + NH 4 NO 3 + 2H 2 O
2NH 4 HS (saturated H 2 S) + 2CuSO 4 = (NH 4) 2 SO 4 + H 2 SO 4 + 2CuS↓
Receipt: saturation of a concentrated solution of NH 3 with hydrogen sulfide:
NH 3 H 2 O (conc.) + H 2 S (g) = NH 4 HS+ H 2 O
In analytical chemistry, a solution containing equal amounts of NH 4 HS and NH 3 H 2 O is conventionally considered a solution of (NH 4) 2 S and the formula of the average salt is used in writing the reaction equations, although ammonium sulfide is completely hydrolyzed in water to NH 4 HS and NH 3H2O.
Sulfur dioxide. Sulfites
Sulfur dioxide SO2. Acidic oxide. Colorless gas with a pungent odor. The molecule has the structure of an incomplete triangle [: S(O) 2 ] (sp 2 - hybridization), contains σ, π bonds S=O. Easily liquefied, thermally stable. Highly soluble in water (~40 l/1 l H 2 O at 20 °C). Forms a polyhydrate with the properties of a weak acid; dissociation products are HSO 3 - and SO 3 2 - ions. The HSO 3 ion has two tautomeric forms - symmetrical(non-acidic) with a tetrahedral structure (sp 3 -hybridization), which predominates in the mixture, and asymmetrical(acidic) with the structure of an incomplete tetrahedron [: S(O) 2 (OH)] (sp 3 -hybridization). The SO 3 2‑ ion is also tetrahedral [: S(O) 3 ].
Reacts with alkalis, ammonia hydrate. A typical reducing agent, weak oxidizing agent.
Qualitative reaction– discoloration of yellow-brown “iodine water”. Intermediate product in the production of sulfites and sulfuric acid.
It is used for bleaching wool, silk and straw, canning and storing fruits, as a disinfectant, antioxidant, and refrigerant. Poisonous.
The compound of composition H 2 SO 3 (sulfurous acid) is unknown (does not exist).
Equations of the most important reactions:
Solubility in water and acidic properties:
Receipt: in industry - combustion of sulfur in air enriched with oxygen, and, to a lesser extent, roasting of sulfide ores (SO 2 - associated gas when roasting pyrite):
S + O 2 = SO 2(280–360 °C)
4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8 SO 2(800 °C, firing)
in the laboratory - displacement of sulfites with sulfuric acid:
BaSO 3 (t) + H 2 SO 4 (conc.) = BaSO 4 ↓ + SO 2 + H 2 O
Sodium sulfite Na 2 SO 3. Oxosol. White. When heated in air, it decomposes without melting and melts under excess pressure of argon. When wet and in solution, it is sensitive to atmospheric oxygen. It is highly soluble in water and hydrolyzes at the anion. Decomposes with acids. Typical reducer.
Qualitative reaction on the SO 3 2‑ ion - the formation of a white precipitate of barium sulfite, which is transferred into solution with strong acids (HCl, HNO 3).
It is used as a reagent in analytical chemistry, a component of photographic solutions, and a chlorine neutralizer for bleaching fabrics.
Equations of the most important reactions:
Receipt:
Na 2 CO 3 (conc.) + SO 2 = Na2SO3+CO2
Sulfuric acid. Sulfates
Sulfuric acid H 2 SO 4. Oxoacid. Colorless liquid, very viscous (oily), very hygroscopic. The molecule has a distorted tetrahedral structure (sp 3 -hybridization), contains covalent σ-bonds S – OH and σπ-bonds S=O. The SO 4 2‑ ion has a regular tetrahedral structure. It has a wide temperature range of the liquid state (~300 degrees). Partially decomposes when heated above 296 °C. It is distilled in the form of an azeotropic mixture with water (mass fraction of acid 98.3%, boiling point 296–340 °C), and with stronger heating it decomposes completely. Unlimitedly miscible with water (with strong exo‑effect). Strong acid in solution, neutralized by alkalis and ammonia hydrate. Converts metals into sulfates (with an excess of concentrated acid under normal conditions, soluble hydrosulfates are formed), but the metals Be, Bi, Co, Fe, Mg and Nb are passivated in concentrated acid and do not react with it. Reacts with basic oxides and hydroxides, decomposes salts of weak acids. A weak oxidizing agent in a dilute solution (due to H I), a strong oxidizing agent in a concentrated solution (due to S VI). It dissolves SO 3 well and reacts with it (a heavy oily liquid is formed - oleum, contains H 2 S 2 O 7).
Qualitative reaction on the SO 4 2‑ ion – precipitation of white barium sulfate BaSO 4 (the precipitate is not transferred into solution by hydrochloric and nitric acids, unlike the white precipitate BaSO 3).
It is used in the production of sulfates and other sulfur compounds, mineral fertilizers, explosives, dyes and drugs, in organic synthesis, for the “opening” (the first stage of processing) of industrially important ores and minerals, in the purification of petroleum products, in the electrolysis of water, as an electrolyte for lead batteries . Toxic, causes skin burns. Equations of the most important reactions:
Receipt V industry:
a) synthesis of SO 2 from sulfur, sulfide ores, hydrogen sulfide and sulfate ores:
S + O 2 (air) = SO 2(280–360 °C)
4FeS 2 + 11O 2 (air) = 8 SO 2+ 2Fe 2 O 3 (800 °C, firing)
2H 2 S + 3O 2 (g) = 2 SO 2+ 2H 2 O (250–300 °C)
CaSO 4 + C (coke) = CaO + SO 2+ CO (1300–1500 °C)
b) conversion of SO 2 to SO 3 in a contact apparatus:
c) synthesis of concentrated and anhydrous sulfuric acid:
H 2 O (dil. H 2 SO 4) + SO 3 = H2SO4(conc., anhydrous)
(the absorption of SO 3 with pure water to produce H 2 SO 4 is not carried out due to the strong heating of the mixture and the reverse decomposition of H 2 SO 4, see above);
d) synthesis oleum– a mixture of anhydrous H 2 SO 4, disulfuric acid H 2 S 2 O 7 and excess SO 3. Dissolved SO 3 guarantees the anhydrity of oleum (when water enters, H 2 SO 4 is immediately formed), which allows it to be safely transported in steel tanks.
Sodium sulfate Na 2 SO 4. Oxosol. White, hygroscopic. Melts and boils without decomposition. Forms crystalline hydrate (mineral mirabilite), easily losing water; technical name Glauber's salt. It is highly soluble in water and does not hydrolyze. Reacts with H 2 SO 4 (conc.), SO 3 . It is reduced by hydrogen and coke when heated. Enters into ion exchange reactions.
It is used in the production of glass, cellulose and mineral paints, as a medicine. Contained in the brine of salt lakes, in particular in the Kara-Bogaz-Gol Bay of the Caspian Sea.
Equations of the most important reactions:
Potassium hydrogen sulfate KHSO 4. Acid oxo salt. White, hygroscopic, but does not form crystalline hydrates. When heated, it melts and decomposes. It is highly soluble in water; the anion undergoes dissociation in solution; the solution environment is strongly acidic. Neutralized by alkalis.
It is used as a component of fluxes in metallurgy, an integral part of mineral fertilizers.
Equations of the most important reactions:
2KHSO 4 = K 2 SO 4 + H 2 SO 4 (up to 240 °C)
2KHSO 4 = K 2 S 2 O 7 + H 2 O (320–340 °C)
KHSO 4 (dil.) + KOH (conc.) = K 2 SO 4 + H 2 O KHSO 4 + KCl = K 2 SO 4 + HCl (450–700 °C)
6KHSO 4 + M 2 O 3 = 2KM(SO 4) 2 + 2K 2 SO 4 + 3H 2 O (350–500 °C, M = Al, Cr)
Receipt: treatment of potassium sulfate with concentrated (more than 6O%) sulfuric acid in the cold:
K 2 SO 4 + H 2 SO 4 (conc.) = 2 KHSO 4
Calcium sulfate CaSO 4. Oxosol. White, very hygroscopic, refractory, decomposes when heated. Natural CaSO 4 occurs as a very common mineral gypsum CaSO 4 2H 2 O. At 130 °C, gypsum loses some of the water and turns into burnt (plaster) gypsum 2CaSO 4 H 2 O (technical name alabaster). Completely dehydrated (200 °C) gypsum corresponds to the mineral anhydrite CaSO4. Slightly soluble in water (0.206 g/100 g H 2 O at 20 °C), solubility decreases when heated. Reacts with H 2 SO 4 (conc.). Restored by coke during fusion. Determines most of the “permanent” hardness of fresh water (see 9.2 for details).
Equations of the most important reactions: 100–128 °C
It is used as a raw material in the production of SO 2, H 2 SO 4 and (NH 4) 2 SO 4, as a flux in metallurgy, and as a paper filler. A binder mortar made from burnt gypsum “sets” faster than a mixture based on Ca(OH) 2 . Hardening is ensured by the binding of water, the formation of gypsum in the form of a stone mass. Burnt gypsum is used to make plaster casts, architectural and decorative forms and products, partition slabs and panels, and stone floors.
Aluminum-potassium sulfate KAl(SO 4) 2. Double oxo salt. White, hygroscopic. Decomposes when heated strongly. Forms crystalline hydrate - potassium alum. Moderately soluble in water, hydrolyzes with aluminum cation. Reacts with alkalis, ammonia hydrate.
It is used as a mordant for dyeing fabrics, a leather tanning agent, a coagulant for purifying fresh water, a component of compositions for sizing paper, and an external hemostatic agent in medicine and cosmetology. It is formed by the joint crystallization of aluminum and potassium sulfates.
Equations of the most important reactions:
Chromium(III) sulfate - potassium KCr(SO 4) 2. Double oxo salt. Red (hydrate dark purple, technical name chromium-potassium alum). When heated, it decomposes without melting. It is highly soluble in water (the gray-blue color of the solution corresponds to aqua complex 3+), hydrolyzes at the chromium(III) cation. Reacts with alkalis, ammonia hydrate. Weak oxidizing and reducing agent. Enters into ion exchange reactions.
Qualitative reactions on the Cr 3+ ion – reduction to Cr 2+ or oxidation to yellow CrO 4 2‑.
It is used as a leather tanning agent, a mordant for dyeing fabrics, and a reagent in photography. It is formed by the joint crystallization of chromium(III) and potassium sulfates. Equations of the most important reactions:
Manganese (II) sulfate MnSO 4 . Oxosol. White, melts and decomposes when heated. Crystalline hydrate MnSO 4 5H 2 O – red-pink, technical name manganese sulfate. It is highly soluble in water; the light pink (almost colorless) color of the solution corresponds to aquacomplex 2+; hydrolyzes at the cation. Reacts with alkalis, ammonia hydrate. Weak reducing agent, reacts with typical (strong) oxidizing agents.
Qualitative reactions on the Mn 2+ ion – commutation with the MnO 4 ion and the disappearance of the violet color of the latter, oxidation of Mn 2+ to MnO 4 and the appearance of a violet color.
It is used for the production of Mn, MnO 2 and other manganese compounds, as a microfertilizer and analytical reagent.
Equations of the most important reactions:
Receipt:
2MnO 2 + 2H 2 SO 4 (conc.) = 2 MnSO4+ O 2 + 2H 2 O (100 °C)
Iron (II) sulfate FeSO 4 . Oxosol. White (light green hydrate, technical name inkstone), hygroscopic. Decomposes when heated. It is highly soluble in water and is slightly hydrolyzed by the cation. It is quickly oxidized in solution by atmospheric oxygen (the solution turns yellow and becomes cloudy). Reacts with oxidizing acids, alkalis, and ammonia hydrate. Typical reducer.
It is used as a component of mineral paints, electrolytes in electroplating, a wood preservative, a fungicide, and a medicine against anemia. In the laboratory it is often taken in the form of a double salt Fe(NH 4) 2 (SO 4) 2 6H 2 O ( Mohr's salt), more resistant to air.
Equations of the most important reactions:
Receipt:
Fe + H 2 SO 4 (diluted) = FeSO4+H2
FeCO 3 + H 2 SO 4 (diluted) = FeSO4+ CO 2 + H 2 O
7.4. Non-metals VA‑group
Nitrogen. Ammonia
Nitrogen– element of the 2nd period and VA group of the Periodic system, serial number 7. Electronic formula of the atom [ 2 He]2s 2 2p 3, characteristic oxidation states 0, ‑III, +III and +V, less often +II, +IV and etc.; the N v state is considered relatively stable.
Scale of nitrogen oxidation states:
Nitrogen has a high electronegativity (3.07), third after F and O. It exhibits typical non-metallic (acidic) properties. Forms various oxygen-containing acids, salts and binary compounds, as well as the ammonium cation NH 4 + and its salts.
In nature - seventeenth by chemical abundance element (ninth among non-metals). A vital element for all organisms.
Nitrogen N 2. Simple substance. It consists of non-polar molecules with a very stable σππ-bond N ≡ N, this explains the chemical inertness of nitrogen under normal conditions. A colorless, tasteless and odorless gas that condenses into a colorless liquid (unlike O2).
Main component of air: 78.09% by volume, 75.52% by mass. Nitrogen boils away from liquid air before oxygen O2. Slightly soluble in water (15.4 ml/1 l H 2 O at 20 ° C), the solubility of nitrogen is less than that of oxygen.
At room temperature, N2 reacts only with lithium (in a humid atmosphere), forming lithium nitride Li3N; nitrides of other elements are synthesized with strong heating:
N 2 + 3Mg = Mg 3 N 2 (800 °C)
In an electrical discharge, N2 reacts with fluorine and, to a very small extent, with oxygen:
The reversible reaction to produce ammonia occurs at 500 °C, under pressure up to 350 atm and always in the presence of a catalyst (Fe/F 2 O 3 /FeO, in the laboratory Pt):
According to Le Chatelier's principle, an increase in ammonia yield should occur with increasing pressure and decreasing temperature. However, the reaction rate at low temperatures is very low, so the process is carried out at 450–500 °C, achieving a 15% ammonia yield. Unreacted N 2 and H 2 are returned to the reactor and thereby increase the degree of reaction.
Nitrogen is chemically passive in relation to acids and alkalis and does not support combustion.
Receipt V industry– fractional distillation of liquid air or removal of oxygen from air by chemical means, for example, by the reaction 2C (coke) + O 2 = 2CO when heated. In these cases, nitrogen is obtained, which also contains impurities of noble gases (mainly argon).
IN laboratories small amounts of chemically pure nitrogen can be obtained by the commutation reaction with moderate heating:
N ‑III H 4 N III O 2(t) = N 2 0 + 2H 2 O (60–70 °C)
NH 4 Cl (p) + KNO 2 (p) = N 2 0 + KCl + 2H 2 O (100 °C)
It is used for the synthesis of ammonia, nitric acid and other nitrogen-containing products, as an inert medium for chemical and metallurgical processes and storage of flammable substances.
Ammonia NH3. Binary compound, the oxidation state of nitrogen is – III. Colorless gas with a sharp characteristic odor. The molecule has the structure of an incomplete tetrahedron [: N(H) 3)] (sp 3 -hybridization). The presence of a donor pair of electrons on the sp 3 -hybrid orbital of nitrogen in the NH 3 molecule determines the characteristic reaction of addition of a hydrogen cation, which results in the formation of a cation ammonium NH4+. It liquefies under excess pressure at room temperature. In the liquid state, it is associated through hydrogen bonds. Thermally unstable. Highly soluble in water (more than 700 l/1 l H 2 O at 20 °C); the proportion in the saturated solution is = 34% by mass and = 99% by volume, pH = 11.8.
Very reactive, prone to addition reactions. Cr reacts in oxygen, reacts with acids. It exhibits reducing (due to N‑III) and oxidizing (due to H I) properties. It is dried only with calcium oxide.
Qualitative reactions– formation of white “smoke” upon contact with gaseous HCl, blackening of a piece of paper moistened with a solution of Hg 2 (NO 3) 2.
An intermediate product in the synthesis of HNO 3 and ammonium salts. Used in the production of soda, nitrogen fertilizers, dyes, explosives; liquid ammonia is a refrigerant. Poisonous.
Equations of the most important reactions:
Receipt: V laboratories– displacement of ammonia from ammonium salts when heated with soda lime (NaOH + CaO):
or boiling an aqueous solution of ammonia and then drying the gas.
IN industry ammonia is synthesized from nitrogen (see) with hydrogen. Produced by industry either in liquefied form or in the form of a concentrated aqueous solution under the technical name ammonia water.
Ammonia hydrate NH 3 H 2 O. Intermolecular connection. White, in the crystal lattice - molecules NH 3 and H 2 O, connected by a weak hydrogen bond H 3 N ... HON. Present in an aqueous solution of ammonia, a weak base (dissociation products - NH 4 ‑ cation and OH ‑ anion). The ammonium cation has a regular tetrahedral structure (sp 3 hybridization). Thermally unstable, completely decomposes when the solution is boiled. Neutralized by strong acids. Shows reducing properties (due to N III) in a concentrated solution. Enters into ion exchange and complexation reactions.
Qualitative reaction– formation of white “smoke” upon contact with gaseous HCl.
It is used to create a slightly alkaline environment in solution during the precipitation of amphoteric hydroxides.
A 1M ammonia solution contains mainly NH 3 H 2 O hydrate and only 0.4% NH 4 + and OH - ions (due to hydrate dissociation); Thus, the ionic “ammonium hydroxide NH 4 OH” is practically not contained in the solution, and there is no such compound in the solid hydrate. Equations of the most important reactions:
NH 3 H 2 O (conc.) = NH 3 + H 2 O (boiling with NaOH)
NH 3 H 2 O + HCl (diluted) = NH 4 Cl + H 2 O
3(NH 3 H 2 O) (conc.) + CrCl 3 = Cr(OH) 3 ↓ + 3NH 4 Cl
8(NH 3 H 2 O) (conc.) + ZBr 2 (p) = N 2 + 6NH 4 Br + 8H 2 O (40–50 °C)
2(NH 3 H 2 O) (conc.) + 2KMnO 4 = N 2 + 2MnO 2 ↓ + 4H 2 O + 2KOH
4(NH 3 H 2 O) (conc.) + Ag2O= 2OH + 3H2O
4(NH 3 H 2 O) (conc.) + Cu(OH) 2 + (OH) 2 + 4H 2 O
6(NH 3 H 2 O) (conc.) + NiCl 2 = Cl 2 + 6H 2 O
A dilute ammonia solution (3–10%) is often called ammonia(the name was invented by alchemists), and the concentrated solution (18.5–25%) - ammonia water(produced by industry).
Related information.
Structure and properties of atoms. Sulfur atoms, like oxygen atoms and all other elements of the main subgroup of group VI of D.I. Mendeleev’s Periodic Table, contain six electrons at the outer energy level, two of which are unpaired electrons. However, compared to oxygen atoms, sulfur atoms have a larger radius and a lower electronegativity value, therefore they exhibit more pronounced reducing properties, forming compounds with oxidation states +2, +4, +6.
In relation to less electronegative elements (hydrogen, metals), sulfur exhibits oxidizing properties and acquires an oxidation state of -2.
Sulfur is a simple substance. Sulfur, like oxygen, is characterized by allotropy. There are many known modifications of sulfur with a cyclic or linear structure of molecules of various compositions.
The most stable modification is known as rhombic sulfur, consisting of S 8 molecules (Fig. 116). Its crystals have the form of octahedra with cut corners. They are lemon yellow and translucent, with a melting point of 112.8 °C. All other modifications transform into this modification at room temperature.
Rice. 116.
Model of a rhombic sulfur molecule
It is known, for example, that during crystallization from a melt, monoclinic sulfur is first obtained (needle-shaped crystals, melting point 119.3 ° C), which then turns into orthorhombic (Fig. 117). When pieces of sulfur are heated in a test tube, it melts, turning into a yellow liquid. At a temperature of about 160 °C, liquid sulfur begins to darken and becomes so thick and viscous that it does not even pour out of the test tube, but with further heating it turns into a highly mobile liquid, but retains the same dark brown color. If you pour it into cold water, it hardens into a transparent rubbery mass. This is plastic sulfur. It can also be obtained in the form of threads. However, after a few days it also turns into rhombic sulfur.
Rice. 117.
Interconversions of allotropic modifications of sulfur
Sulfur does not dissolve in water. Sulfur crystals sink in water, but the powder floats on the surface of the water, since small sulfur crystals are not wetted by water and are kept afloat by small air bubbles. This is a flotation process. Sulfur is slightly soluble in ethyl alcohol and diethyl ether, and readily dissolves in carbon disulfide.
Under normal conditions, sulfur reacts with all alkali and alkaline earth metals, copper, mercury, silver, for example:
This reaction underlies the removal and neutralization of spilled mercury, for example from a broken thermometer. Visible drops of mercury can be collected on a sheet of paper or on a copper plate. Mercury that gets into the cracks must be covered with sulfur powder. This process is called demercurization.
When heated, sulfur also reacts with other metals (Zn, Al, Fe). Only gold does not interact with it under any circumstances.
Sulfur also exhibits oxidizing properties with hydrogen, with which it reacts when heated:
H 2 + S = H 2 S.
Of the nonmetals, only nitrogen and iodine, as well as noble gases, do not react with sulfur.
Sulfur burns with a bluish flame, and sulfur oxide (IV) is formed:
S + O 2 = SO 2.
This compound is commonly known as sulfur dioxide.
Laboratory experiment No. 28
Combustion of sulfur in air and oxygen
In nature, sulfur occurs in three forms: native, sulfide and sulfate (Fig. 118, Table 8).
Rice. 118.
Sulfur in nature:
1 - native sulfur; 2 - pyrite; 3 - zinc blende; 4 - gypsum; 5 - Glauber's salt
Table 8
Sulfur in nature
Sulfur is a vital chemical element. It is part of proteins - one of the main chemical components of the cells of all living organisms. There is especially a lot of sulfur in the proteins of hair, horns, and wool. In addition, sulfur is an integral part of biologically active substances in the body: vitamins and hormones (for example, insulin).
Sulfur is involved in the redox processes of the body. With a lack of sulfur in the body, fragility and brittleness of bones and hair loss occur.
Legumes (peas, lentils), oatmeal, and eggs are rich in sulfur.
Application of sulfur. Sulfur has been known to people since ancient times. It gets its name from the Sanskrit word sira, which means “light yellow.” Sulfur was used in Ancient Egypt already two thousand years BC for the preparation of paints, bleaching fabrics and making cosmetics. In Ancient Rome, sulfur was used to treat skin diseases, and in Ancient Greece it was burned to disinfect things and indoor air.
In the Middle Ages, among alchemists, sulfur was an expression of one of the “fundamental principles of nature” and an obligatory component of the “philosopher’s stone”.
If you have read the famous novel by A. Dumas “The Count of Monte Cristo”, you will be able to name the areas of application of sulfur that have been known since ancient times. The hero of the novel, Abbot Faria, pretended to have a skin disease, and to treat it he was given sulfur, which the enterprising abbot used to make gunpowder.
Sulfur is used in the production of matches and paper, rubber and paints, explosives and medicines, cosmetics. In agriculture, it is used to combat pathogens of fungal and bacterial diseases, and plant pests (Fig. 119).
Rice. 119.
Application of sulfur:
1 - production of ointments; 2 - production of matches; 3 - production of explosives; 4 - production of sulfuric acid; 5 - pulp and paper industry; 6 - in agriculture for disinfection of premises; 7 - getting rubber
However, the main consumer of sulfur is the chemical industry. About half of the world's sulfur is used to produce sulfuric acid.
New words and concepts
- The structure of sulfur atoms and the oxidation state of sulfur.
- Allotropy of sulfur: orthorhombic, monoclinic and plastic sulfur.
- Chemical properties of sulfur: interaction with metals, oxygen, hydrogen. Demercurization.
- Sulfur in nature: native, sulfide and sulfate sulfur.
- Biological significance of sulfur.
- Application of sulfur.
