Sulfur. Hydrogen sulfide and sulfides
4.doc
240Sulfur. Hydrogen sulfide, sulfides, hydrosulfides. Sulfur oxides (IV) and (VI). Sulfuric acid and sulfuric acid and their salts. Esters of sulfuric acid. Sodium Thiosulfate
4.1. Sulfur
Sulfur is one of the few chemical elements that people have used for several millennia. It is widely distributed in nature and is found both in the free state (native sulfur) and in the compounds. Minerals containing sulfur can be divided into two groups - sulfides (pyrites, glitters, dummies) and sulfates. Native sulfur is found in large quantities in Italy (Sicily Island) and the USA. In the CIS, native sulfur deposits are found in the Volga region, in the states of Central Asia, in the Crimea and other areas.
The minerals of the first group include lead shine PbS, copper shine Cu 2 S, silver shine - Ag 2 S, zinc blende - ZnS, cadmium snag - CdS, pyrite or iron pyrite - FeS 2, chalcopyrite - CuFeS 2, cinnabar - HgS.
The minerals of the second group include the gypsum CaSO 4 2H 2 O, mirabilite (Glauber's salt) - Na 2 SO 4 10H 2 O, and kizerite - MgSO 4 H 2 O.
Sulfur is found in animals and plants, as part of the protein molecules. Organic sulfur compounds are found in petroleum.
Getting
1. When sulfur is obtained from natural compounds, for example, from pyritic sulfur, it is heated to high temperatures. Sulfur pyrite decomposes to form iron (II) sulfide and sulfur:
2. Sulfur can be obtained by oxidation of hydrogen sulfide by the lack of oxygen by the reaction:
2H 2 S O 2 = 2S 2H 2 O
3. Nowadays, sulfur recovery by carbon dioxide reduction of sulfur dioxide SO 2 is a common by-product in the smelting of metals from sulfur ores:
SO 2 C = CO 2 S
4. The waste gases of metallurgical and coke ovens contain a mixture of sulfur dioxide and hydrogen sulfide. This mixture is passed at a high temperature over the catalyst:
H 2 S SO 2 = 2H 2 O 3S
^ Physical properties
Sulfur is a hard, brittle lemon yellow color. It is practically insoluble in water, but well soluble in CS 2 aniline in carbon disulfide and some other solvents.
Poor conducts heat and electrical current. Sulfur forms several allotropic modifications:
1 . ^ Rhombic sulfur (most stable), crystals have the form of octahedra.
When sulfur is heated, its color and viscosity change: first, light yellow is formed, and then as the temperature increases, it darkens and becomes so viscous that it does not flow out of the tube, with further heating, the viscosity drops again, and at 444, 6 ° С sulfur boils .
2. ^ Monoclinic sulfur - modification in the form of dark yellow needles, obtained by slowly cooling the molten sulfur.
3. Plastic sulfurformed if the sulfur heated to boiling is poured into cold water. Easily stretched like rubber (see fig. 19).
Natural sulfur consists of a mixture of four stable isotopes: 32 16 S, 33 16 S, 34 16 S, 36 16 S.
^ Chemical properties
A sulfur atom, having an incomplete external energy level, can attach two electrons and exhibit a degree
Oxidation -2. Sulfur exhibits such oxidation state in compounds with metals and hydrogen (Na 2 S, H 2 S). When recoil or depletion of electrons to the atom of the more electronegative element, the degree of sulfur oxidation can be 2, 4, 6
Sulfur is relatively inert in the cold, but with increasing temperature, its reactivity increases. 1. With metals, sulfur exhibits oxidizing properties. When these reactions form sulfides (with gold, platinum and iridium does not react): Fe S = FeS
2. Under normal conditions, sulfur does not react with hydrogen, and at 150–200 ° C a reversible reaction occurs:
3. In reactions with metals and with hydrogen, sulfur behaves as a typical oxidizing agent, and in the presence of strong oxidizing agents exhibits reducing properties.
S 3F 2 = SF 6 (does not react with iodine)
4. Combustion of sulfur in oxygen proceeds at 280 ° C, and in air at 360 ° C. This forms a mixture of SO 2 and SO 3:
S O 2 = SO 2 2S 3O 2 = 2SO 3
5. When heated without air, sulfur directly combines with phosphorus, carbon, showing oxidizing properties:
2P 3S = P 2 S 3 2S C = CS 2
6. When interacting with complex substances, sulfur behaves mainly as a reducing agent:
7. Sulfur is capable of disproportionation reactions. So, when boiling sulfur powder with alkalis, sulfites and sulfides are formed:
Application
Sulfur is widely used in industry and agriculture. About half of its production is consumed to produce sulfuric acid. Sulfur is used for the vulcanization of rubber: in this case, the rubber is converted into rubber.
In the form of sulfur (fine powder) sulfur is used to combat diseases of the vineyard and cotton. It is used to obtain gunpowder, matches, luminous compositions. In medicine, prepared sulfur ointment for the treatment of skin diseases.
4.2. Hydrogen sulfide, sulfides, hydrosulfides
Hydrogen sulfide is an analogue of water. Its electronic formula
It shows that two p-electrons of the external level of the sulfur atom are involved in the formation of H – S – H bonds. The H 2 S molecule has an angular shape, so it is polar.
^ Being in nature
Hydrogen sulfide is found in nature in volcanic gases and in the waters of some mineral sources, such as Pyatigorsk, Matsesta. It is formed by the decay of sulfur-containing organic substances from various animal and plant residues. This explains the characteristic unpleasant smell of sewage, cesspools and landfills.
Getting
1. Hydrogen sulfide can be obtained by direct connection of sulfur with hydrogen by heating:
2. But usually it is obtained by the action of dilute hydrochloric or sulfuric acid on iron (III) sulfide:
2HCl FeS = FeCl 2 H 2 S 2H FeS = Fe 2 H 2 S This reaction is often carried out in the Kipp apparatus.
^ Physical properties
Under normal conditions, hydrogen sulfide is a colorless gas with a strong characteristic smell of rotten eggs. Very poisonous, inhalation binds to hemoglobin, causing paralysis, which is unhealthy.
Ko leads to death. In low concentrations, less dangerous. It is necessary to work with it in exhaust cabinets or with hermetically closed devices. The permissible content of H 2 S in the production premises is 0.01 mg per 1 liter of air.
Hydrogen sulfide is relatively well soluble in water (at 20 ° С 2.5 volumes of hydrogen sulfide dissolve in 1 volume of water).
A solution of hydrogen sulfide in water is called hydrogen sulfide water or hydrogen sulfide acid (it detects the properties of a weak acid).
^ Chemical properties
1, When strongly heated, hydrogen sulfide almost completely decomposes to form sulfur and hydrogen.
2. Gaseous hydrogen sulfide burns in the air with a blue flame with the formation of sulfur oxide (IV) and water:
2H 2 S 3O 2 = 2SO 2 2H 2 O
With a lack of oxygen, sulfur and water are formed: 2H 2 S O 2 = 2S 2H 2 O
3. Hydrogen sulfide is a fairly strong reducing agent. This important chemical property can be explained as follows. In a solution of H 2 S, it is relatively easy to donate electrons to air oxygen molecules:
At the same time, oxygen in the air oxidizes hydrogen sulfide to sulfur, which makes the hydrogen sulfide water turbid:
2H 2 S O 2 = 2S 2H 2 O
This explains the fact that hydrogen sulfide does not accumulate in very large quantities in nature when organic matter decays - oxygen from the air oxidizes it to free sulfur.
4, Hydrogen sulphide reacts vigorously with halogen solutions, for example:
H 2 S I 2 = 2HI S Sulfur is liberated and the iodine solution becomes discolored.
5. Various oxidizers react vigorously with hydrogen sulfide: under the action of nitric acid, free sulfur is formed.
6. A solution of hydrogen sulfide is acidic due to dissociation:
H 2 SH HS - HS - H S -2
The first stage usually prevails. It is a very weak acid: weaker than coal acid, which usually displaces H 2 S from sulphides.
Sulphides and hydrosulphides
Hydrofluoric acid, as a dibasic, forms two rows of salts:
Middle - sulfides (Na 2 S);
Sour - hydrosulfides (NaHS).
These salts can be obtained: by the interaction of hydroxides with hydrogen sulfide: 2NaOHH 2 S = Na 2 S 2H 2 O
The direct interaction of sulfur with metals:
The exchange reaction of salts with H 2 S or between salts:
Pb (NO 3) 2 Na 2 S = PbS 2NaNO 3
CuSO 4 H 2 S = CuS H 2 SO 4 Cu 2 H 2 S = CuS 2H
Hydrosulphides are almost all soluble in water.
Sulfides of alkali and alkaline earth metals are also easily soluble in water, colorless.
Heavy metal sulphides are practically insoluble or slightly soluble in water (FeS, MnS, ZnS); Some of them do not dissolve in dilute acids (CuS, PbS, HgS).
As salts of a weak acid, sulfides in aqueous solutions are highly hydrolyzed. For example, alkali metal sulphides, when dissolved in water, have an alkaline reaction:
Na 2 S НОНNaHS NaOH
All sulphides, like hydrogen sulfide itself, are vigorous reducing agents:
3PbS -2 8HN 5 O 3 (par.) = 3PbS 6 O 4 4H 2 O 8N 2 O
Some sulfides have a characteristic color: CuS and PbS - black, CdS - yellow, ZnS - white, MnS - pink, SnS - brown, Al 2 S 3 - orange. The qualitative analysis of cations is based on the different solubilities of sulphides and the different colors of many of them.
^ 4.3. Sulfur Oxide (IV) and Sulfuric Acid
Sulfur oxide (IV), or sulfur dioxide, under normal conditions, is a colorless gas with a sharp suffocating odor. When cooled to -10 ° C, it liquefies into a colorless liquid.
Getting
1. Under laboratory conditions, sulfur oxide (IV) is obtained from salts of sulfurous acid by the action of strong acids on them:
Na 2 SO 3 H 2 SO 4 = Na 2 SO 4 S0 2 H 2 O 2NaHSO 3 H 2 SO 4 = Na 2 SO 4 2SO 2 2H 2 O 2HSO - 3 2H = 2 SO 2 2H 2 O
2. Also sulfur dioxide is formed by the interaction of concentrated sulfuric acid when heated with low-active metals:
Cu 2H 2 SO 4 = CuSO 4 SO 2 2H 2 O
Cu 4H 2 SO 2-4 = Cu 2 SO 2-4 SO 2 2H 2 O
3. Sulfur oxide (IV) is also formed when sulfur is burned in air or oxygen:
4. In industrial conditions, SO 2 is obtained by roasting pyrite with FeS 2 or sulfur ores of non-ferrous metals (zinc blende ZnS, lead gloss PbS, etc.):
4FeS 2 11О 2 = 2Fe 2 O 3 8SO 2
The structural formula of the molecule SO 2:
Four sulfur atoms and four electrons from two oxygen atoms take part in the formation of bonds in the SO 2 molecule. The mutual repulsion of the bonding electron pairs and the lone electron pair of sulfur gives the molecule an angular shape.
Chemical properties
1. Sulfur oxide (IV) exhibits all the properties of acid oxides:
Water interaction
Interaction with alkalis
Interaction with basic oxides.
2. Sulfur oxide (IV) is characterized by reducing properties:
S 4 O 2 O 0 2 2S 6 O -2 3 (in the presence of a catalyst, when heated)
But in the presence of strong reducing agents, SO 2 behaves as an oxidizing agent:
The redox duality of sulfur (IV) oxide is explained by the fact that sulfur has an oxidation state of 4 in it, and therefore it can, giving up 2 electrons, be oxidized to S 6, and by taking 4 electrons, it can be reduced to S °. The manifestation of these or other properties depends on the nature of the reactive component.
Sulfur oxide (IV) is highly soluble in water (40 volumes of SO 2 dissolve in 1 volume at 20 ° С). In this case, sulfurous acid existing only in aqueous solution is formed:
SO 2 H 2 ОH 2 SO 3
The reaction is reversible. In aqueous solution, sulfur oxide (IV) and sulfurous acid are in chemical equilibrium, which can be shifted. With the binding of H 2 SO 3 (neutralization of acid
You) the reaction proceeds in the direction of the formation of sulfurous acid; when removing SO 2 (blowing through a solution of nitrogen or heating), the reaction proceeds towards the starting materials. Sulfuric acid solution always contains sulfur oxide (IV), which gives it a sharp odor.
Sulfuric acid has all the properties of acids. In solution it dissociates in steps:
H 2 SO 3 H HSO - 3 HSO - 3 H SO 2- 3
Thermally unstable, volatile. Sulfurous acid, as dibasic, forms two types of salts:
Middle - sulfites (Na 2 SO 3);
Sour - hydrosulfite (NaHSO 3).
Sulfites are formed when alkali is completely neutralized by the acid:
H 2 SO 3 2NaOH = Na 2 SO 3 2H 2 O
Hydrosulphites are obtained with a lack of alkali:
H 2 SO 3 NaOH = NaHSO 3 H 2 O
Sulfuric acid and its salts possess both oxidizing and reducing properties, which is determined by the nature of the reaction partner.
1. So, under the action of oxygen sulfites are oxidized to sulfates:
2Na 2 S 4 O 3 O 0 2 = 2Na 2 S 6 O -2 4
Sulfurous acid oxidation by bromine and potassium permanganate proceeds even more easily:
5H 2 S 4 O 3 2KMn 7 O 4 = 2H 2 S 6 O 4 2Mn 2 S 6 O 4 K 2 S 6 O 4 3H 2 O
2. In the presence of more vigorous reducing agents, sulfites exhibit oxidative properties:
Almost all hydrosulphites and alkali metal sulphites dissolve from salts of sulfurous acid.
3. Since H 2 SO 3 is a weak acid, the action of acids on sulfites and hydrosulfites results in the release of SO 2. This method is usually used when obtaining SO 2 in laboratory conditions:
NaHSO 3 H 2 SO 4 = Na 2 SO 4 SO 2 H 2 O
4. Water-soluble sulfites readily undergo hydrolysis, as a result of which the concentration of OH - ions increases in the solution:
Na 2 SO 3 NONNaHSO 3 NaOH
Application
Sulfur oxide (IV) and sulfurous acid discolor many dyes, forming with them colorless compounds. The latter can decompose again when heated or in the light, as a result of which the color is restored. Therefore, the whitening effect of SO 2 and H 2 SO 3 is different from the whitening effect of chlorine. Usually, sulfur (IV) rxid bleaches wool, silk and straw.
Sulfur oxide (IV) kills many microorganisms. Therefore, to destroy mold fungi, they fumigate raw cellars, cellars, wine barrels, etc. It is also used in the transport and storage of fruits and berries. In large quantities, sulfur oxide IV) is used to produce sulfuric acid.
An important application is the solution of calcium hydrosulfite CaHSO 3 (sulphite liquor), which is used to treat wood and paper pulp.
^ 4.4. Sulfur oxide (VI). Sulphuric acid
Sulfur oxide (VI) (see tab. 20) is a colorless liquid that solidifies at a temperature of 16.8 ° С to a solid crystalline mass. It absorbs moisture very strongly, forming sulfuric acid: SO 3 H 2 O = H 2 SO 4
Table 20. Properties of sulfur oxides
The dissolution of sulfur oxides (VI) in water is accompanied by the release of a significant amount of heat.
Sulfur oxide (VI) is very soluble in concentrated sulfuric acid. A solution of SO 3 in an anhydrous acid is called oleum. Oleums can contain up to 70% SO 3.
Getting
1. Sulfur oxide (VI) is obtained by oxidation of sulfur dioxide with oxygen in the presence of catalysts at a temperature of 450 ° С (see Sulfuric acid production):
2SO 2 O 2 = 2SO 3
2. Another way to oxidize SO 2 to SO 3 is to use nitric (IV) oxide as an oxidizing agent:
The resulting nitric oxide (II) when interacting with atmospheric oxygen easily and quickly turns into nitric oxide (IV): 2NO O 2 = 2NO 2
Which can again be used in the oxidation of SO 2. Therefore, NO 2 acts as an oxygen carrier. This method of oxidizing SO 2 to SO 3 is called nitrous. The molecule of SO 3 has the shape of a triangle, in the center of which
The sulfur atom is located:
Such a structure is due to the mutual repulsion of the bonding electron pairs. A sulfur atom has provided six external electrons for their formation.
Chemical properties
1. SO 3 - typical acid oxide.
2. Sulfur oxide (VI) has the properties of a strong oxidizing agent.
Application
Sulfur oxide (VI) is used to produce sulfuric acid. The most important is the contact method of obtaining
Sulfuric acid. By this method, you can get H 2 SO 4 of any concentration, as well as oleum. The process consists of three stages: obtaining SO 2; oxidation of SO 2 to SO 3; getting H 2 SO 4.
SO 2 is obtained by burning pyrite FeS 2 in special furnaces: 4FeS 2 11О 2 = 2Fe 2 O 3 8SO 2
To accelerate roasting, pyrite is pre-ground, and for more complete burning out of sulfur, significantly more air (oxygen) is injected than is required by the reaction. The gas leaving the kiln is made up of sulfur (IV) oxide, oxygen, nitrogen, arsenic compounds (from impurities in pyrites) and water vapor. It is called roasting gas.
The roasting gas undergoes thorough cleaning, since even a small content of arsenic compounds, as well as dust and moisture, poison the catalyst. Gas is cleaned from arsenic compounds and dust, passing it through special electrostatic precipitators and a washing tower; moisture is absorbed by concentrated sulfuric acid in a drying tower. The purified gas containing oxygen is heated in a heat exchanger up to 450 ° C and enters the contact apparatus. Inside the contact apparatus are lattice shelves filled with catalyst.
Previously, finely crushed metallic platinum was used as a catalyst. Subsequently, it was replaced by vanadium compounds - vanadium oxide (V) V 2 O 5 or vanadyl sulfate VOSO 4, which is cheaper than platinum and more slowly poisoned.
The reaction of oxidation of SO 2 to SO 3 is reversible:
2SO 2 O 2 2SO 3
An increase in the oxygen content in the calcining gas increases the yield of sulfur oxide (VI): at a temperature of 450 ° C, it usually reaches 95% or more.
The formed sulfur oxide (VI) is then fed by a counter-current method to an absorption tower, where it is absorbed by concentrated sulfuric acid. As it saturates, anhydrous sulfuric acid first forms, and then oleum. In the future, oleum is diluted to 98% sulfuric acid and delivered to consumers.
The structural formula of sulfuric acid:
^ Physical properties
Sulfuric acid is a heavy, colorless, oily liquid that crystallizes at 10.4 ° C, almost double ( = 1.83 g / cm 3) is heavier than water, odorless, nonvolatile. Extremely hygroscopic. Moisture absorbs with the release of large amounts of heat, so you can not pour water to the concentrated sulfuric acid - there will be a splashing of acid. For the
Sulfuric acid must be added to water in small portions.
Anhydrous sulfuric acid dissolves up to 70% of sulfur oxide (VI). When heated, it removes SO 3 until a solution with a mass fraction of H 2 SO 4 of 98.3% is formed. Anhydrous H 2 SO 4 almost does not conduct electric current.
^ Chemical properties
1. It mixes with water in any proportions and forms hydrates of different composition:
H 2 SO 4 H 2 O, H 2 SO 4 2H 2 O, H 2 SO 4 3H 2 O, H 2 SO 4 4H 2 O, H 2 SO 4 6.5H 2 O
2. Concentrated sulfuric acid carbonizes organic matter - sugar, paper, wood, fiber, taking away from them the elements of water:
C 12 H 22 O 11 H 2 SO 4 = 12 C H 2 SO 4 11H 2 O
The formed coal partially interacts with the acid:
Gas dehydration is based on water absorption by sulfuric acid.
As a strong non-volatile acid, H 2 SO 4 displaces other acids from dry salts:
NaNO 3 H 2 SO 4 = NaHSO 4 HNO 3
However, if H 2 SO 4 is added to salt solutions, no acid displacement occurs.
H 2 SO 4 - strong dibasic acid: H 2 SO 4 H HSO - 4 HSO - 4 H SO 2- 4
It has all the properties of non-volatile strong acids.
Diluted sulfuric acid is characterized by all the properties of non-oxidizing acids. Namely: it interacts with metals that are in the electrochemical series of the voltage of metals up to hydrogen:
Interaction with metals is due to the reduction of hydrogen ions.
6. Concentrated sulfuric acid is an energetic oxidizing agent. When heated, oxidizes most metals, including those in the electrochemical series of stresses after hydrogen, does not react only with platinum and gold. Depending on the activity of the metal, the reduction products can be S -2, S ° and S 4.
In the cold, concentrated sulfuric acid does not interact with such strong metals as aluminum, iron, and chromium. This is due to the passivation of metals. This feature is widely used when it is transported in an iron container.
However, when heated:
Thus, concentrated sulfuric acid interacts with metals due to the reduction of the acid-forming atoms.
A qualitative reaction to the sulfate ion SO 2-4 is the formation of a white crystalline precipitate of BaSO 4, insoluble in water and acids:
SO 2- 4 Ba 2 BaSO 4
Application
Sulfuric acid is the most important product of the main chemical industry engaged in the production of non-
Organic acids, alkalis, salts, mineral fertilizers and chlorine.
For a variety of applications sulfuric acid ranks first among the acids. The greatest amount of it is consumed to obtain phosphate and nitrogen fertilizers. Being non-volatile, sulfuric acid is used to produce other acids - hydrochloric, hydrofluoric, phosphoric and acetic.
A lot of it goes for the purification of petroleum products - gasoline, kerosene, lubricating oils - from harmful impurities. In mechanical engineering, sulfuric acid cleans the metal surface from oxides before coating (nickel plating, chrome plating, etc.). Sulfuric acid is used in the manufacture of explosives, artificial fibers, dyes, plastics and many others. It is used to fill the batteries.
Sulfuric acid salts are important.
^ Sodium SulphateNa 2 SO 4 crystallizes from aqueous solutions in the form of a hydrate of Na 2 SO 4 10H 2 O, which is called the Glauber's salt. It is used in medicine as a laxative. Anhydrous sodium sulfate is used in the manufacture of soda and glass.
^ Ammonium Sulphate(NH 4) 2 SO 4 - nitrogen fertilizer.
Potassium sulfateK 2 SO 4 - potash fertilizer.
Calcium sulfate CaSO 4 is found in nature in the form of the gypsum mineral CaSO 4 2H 2 O. When heated to 150 ° C, it loses some of the water and goes into a 2CaSO 4 H 2 O hydrate, called burnt gypsum, or alabaster. When mixed with water into a pasty mass, alabaster hardens again after some time, turning into CaSO 4 2H 2 O. Gypsum is widely used in construction (plaster).
Magnesium sulfateMgSO 4 is found in seawater, causing its bitter taste. Crystalline hydrate, called the bitter salt, is used as a laxative.
Vitriol- technical name of crystalline hydrates of metal sulfates Fe, Cu, Zn, Ni, Co (dehydrated salts are not vitrials). Copper sulfateCuSO 4 5H 2 O is a blue toxic substance. Plants are sprayed with a diluted solution and seeds are sown before sowing. inkstoneFeSO 4 7H 2 O is a light green substance. Used to control pests of plants, inks, mineral paints, etc. Zinc sulfateZnSO 4 7H 2 O is used in the production of mineral inks, in sittoprechatanii, medicine.
^ 4.5. Esters of sulfuric acid. Sodium thiosulfate
Sulfuric acid esters include dialkyl sulfates (RO 2) SO 2. These are high boiling liquids; lower soluble in water; in the presence of alkali, alcohol and salts of sulfuric acid are formed. Lower dialkyl sulfates are alkylating agents.
Diethyl sulfate(C 2 H 5) 2 SO 4. Melting point -26 ° С, boiling point 210 ° С, soluble in alcohols, insoluble in water. Obtained by reacting sulfuric acid with ethanol. Is the ethylating agent in organic synthesis. Penetrates skin.
Dimethyl sulfate(CH 3) 2 SO 4. Melting point -26.8 ° С, boiling point 188.5 ° С. Soluble in alcohols, bad - in water. Reacts with ammonia in the absence of solvent (with an explosion); Sulfur some aromatic compounds, such as phenol esters. It is obtained by interaction of 60% oleum with methanol at 150 ° С. It is a methylating agent in organic synthesis. Carcinogenic, affects the eyes, skin, respiratory organs.
^ Sodium thiosulfate Na 2 S 2 O 3
Salt of thiosulfuric acid in which two sulfur atoms have different oxidation states: 6 and -2. Crystalline substance, well soluble in water. Available in the form of crystalline Na 2 S 2 O 3 5H 2 O, commonly referred to as hyposulphite. Obtained by the interaction of sodium sulfite with sulfur during boiling:
Na 2 SO 3 S = Na 2 S 2 O 3
Like thiosulfuric acid, it is a strong reducing agent. It is easily oxidized by chlorine to sulfuric acid:
Na 2 S 2 O 3 4Cl 2 5H 2 O = 2H 2 SO 4 2NaCl 6HCl
The use of sodium thiosulfate for chlorine absorption (in the first gas masks) was based on this reaction.
Slightly different oxidation occurs sodium thiosulfate weak oxidizing agents. This forms salts of tetrathionic acid, for example:
2Na 2 S 2 O 3 I 2 = Na 2 S 4 O 6 2NaI
Sodium thiosulfate is a by-product in the production of NaHSO 3, sulfur dyes, in the purification of industrial gases from sulfur. It is used to remove traces of chlorine after bleaching fabrics, to extract silver from ores; is a fixative in photography, a reagent in iodometry, an antidote for poisoning with arsenic compounds, mercury, an anti-inflammatory agent.
The chemical element sulfur is located in the VIА-group, period 3 PSCE. On the outer electron layer of the sulfur atom there are 6 electrons. Showing the properties of the oxidizing agent in the formation of substances with metals and hydrogen, sulfur acquires the oxidation state -2 (takes 2 electrons).
In the composition of oxygen-containing compounds, sulfur exhibits oxidation states +4 and +6. Thus, sulfur is characterized by oxidation states -2; 0; +4; +6.
In nature, sulfur is found in its native state and in the composition of minerals. For example, sulfur contains lead gloss (the main component is lead sulfide PbS) and copper gloss (the main component is copper sulfide Cu 2 S).
Sulfur forms a few simple substances - allotropic modifications. Rhombic sulfur is most stable at room temperature (Fig. 1). This substance consists of S 8 molecules. Rhombic sulfur is yellow and melts at a temperature of + 112.8 ° C.
Fig. 1. Rhombic modification of sulfur
When heated, rhombic sulfur gradually turns into a viscous dark brown mass. This is another allotropic modification of sulfur - plastic sulfur. Plastic sulfur consists of linear sulfur molecules S n.
Sulfur in chemical reactions with metals and hydrogen plays the role of an oxidizing agent. Its degree of oxidation decreases from 0 to -2. In the reaction with oxygen, sulfur acts as a reducing agent, increasing its degree of oxidation from 0 to +4.
Consider examples of interactions involving sulfur.
When sulfur interacts with hydrogen, hydrogen sulfide is formed:
When zinc interacts with sulfur, zinc sulfide is formed:
Sulfur burns in oxygen with the formation of sulfur oxide (IV) (Fig. 2):
Fig. 2. Burning sulfur in oxygen
Consider the properties of sulfur compounds with oxidation state "-2". Such compounds include hydrogen sulfide and sulfides - salts of hydrogen sulfide acid.
Hydrogen sulfide is a gas with the smell of rotten eggs. It burns in the air. Moreover, with a lack and an excess of oxygen, combustion proceeds differently.
In an excess of oxygen, hydrogen sulfide burns to form sulfur oxide (IV) and water:
2H 2 S + 3O 2 = 2SO 2 + 2H 2 O.
With a lack of oxygen, incomplete combustion of hydrogen sulfide occurs with the release of sulfur: 2H 2 S + O 2 = 2S + 2H 2 O.
Hydrogen sulfide is highly soluble in water. The resulting solution is a weak hydrogen sulfide acid. Hydrogen sulfide salts are called sulfides. Hydrofluoric acid and water-soluble sulphides enter into exchange reactions.
The interaction of hydrogen sulfide acid and copper (II) chloride forms insoluble copper (II) sulfide and hydrochloric acid: H 2 S + CuCl 2 = CuS + 2HCl.
When potassium sulfide interacts with zinc nitrate, zinc sulfide precipitates and potassium nitrate is formed: K 2 S + Zn (NO 3) 2 = ZnS + 2 KNO 3.
Bibliography
- Orzhekovsky P.A. Collection of tasks and exercises in chemistry: 9th grade: to the textbook P.A. Orzhekovsky and others. “Chemistry. Grade 9 / P.A. Orzhekovsky, N.A. Titov, F.F. Hegel - M .: AST: Astrel, 2007. (p. 91-97)
- Orzhekovsky P.A. Chemistry: 9th grade: studies. for general image. institution / P.A. Orzhekovsky, L.M. Meshcheryakova, L.S. Pontak - M .: AST: Astrel, 2007. (§ 34)
- Orzhekovsky P.A. Chemistry: 9th grade: studies for general. institution / P.A. Orzhekovsky, L.M. Meshcheryakova, M.M. Shalashov. - M .: Astrel, 2013. (§§ 20, 21)
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Homework
- with. 136 No. 3; p.140 №№ 2-4 from the textbook P.A. Orzhekovsky "Chemistry: 9th grade" / P.A. Orzhekovsky, L.M. Meshcheryakova, M.M. Shalashov. - M .: Astrel, 2013.
Sulphides (from lat. sulphur, sulfur - sulfur)
sulfur compounds with more electropositive elements; can be considered as salts of hydrogen sulfide acid (See Hydrogen sulfide acid) H 2 S. There are two C series: the average (normal) of the general formula M 2 S and the acid (hydrosulfides) of the general formula MHS, where M is a monovalent metal. C. alkali metals are colorless, soluble in water. Their aqueous solutions are highly hydrolyzed and alkaline. Under the action of dilute acids emit H 2 S. C. alkaline earth metals are colorless, slightly soluble in water. H 2 S is emitted in humid air. Other properties are similar to S. alkali metals. Both those and others S. are easily oxidized to sulphates. C. heavy metals are practically insoluble in water. Almost all of them are black or black-brown (with the exception of white ZnS, pinkish MnS, yellow CdS, orange-red Sb 2 S 3, yellow SnS 2). The unequal ratio of C. to acids and C. ammonium is used in chemical analysis. I.K. Malina.
Great Soviet Encyclopedia. - M .: Soviet Encyclopedia. 1969-1978 .
See what "Sulphides" are in other dictionaries:
Sulfur compounds with metals and some non-metals. Metal sulfides are salts of hydrogen sulfide acid H2S: medium (for example, Na2S) and acidic, or hydrosulfides (NaHS). By burning natural sulphides, non-ferrous metals and SO2 are obtained. Phosphor sulfides and ... ... Big Encyclopedic Dictionary
Modern Encyclopedia
Sulphides - SULFIDES, inorganic sulfides of sulfur compounds with metals and some non-metals. Included in sulphide ores; used as phosphors (for example, CdS, ZnS). Molybdenum sulfide, titanium solid lubricants. Phosphorus sulphides ... ... Illustrated Encyclopedic Dictionary
- (new lat., from the Latin. sulfur sulfur). Compounds of some body with sulfur, corresponding to oxides or acids. Dictionary of foreign words included in the Russian language. Chudinov, AN, 1910. SULFIDES Novolatinsk., From Lat. sulfur, sulfur. Connection ... ... Dictionary of foreign words of the Russian language
1. Natural sulfur compounds of metals and some non-metals. In the chemical relation are considered as salts of hydrogen sulfide acid H2S. A number of elements form polysulfides with sulfur, which are salts of the poly sulphurous acid H2Sx. The main ... ... Geological encyclopedia
Sulphides - - sulfur compounds with metals and some non-metals; metal sulfides salts of hydrogen sulfide acid H2S. [Terminological dictionary for concrete and reinforced concrete. FSUE “SIC“ Construction ”NIIZHB and A. A. Gvozdeva metro station, Moscow, 2007. 110 p.] ...… Encyclopedia of terms, definitions and explanations of building materials
SULPHES, s, u sulfide, a, husband (specialist.). Chemical compounds of sulfur with metals and certain non-metals. Organic with Natural s. | adj sulphide, oh, oh. Dictionary Ozhegova. S.I. Ozhegov, N.Yu. Shvedov. 1949 1992 ... Dictionary Ozhegova
Sulfides, R2S (R aromatic radical), are most easily obtained by adding dropwise a solution of diazo salts to 60 70 ... Encyclopedia of Brockhaus and Efron
SULFIDES - (1) in inorganic chemistry, such compounds of elements with sulfur, in which the sulfur atoms have an oxidation state of 2. In chemical. relation are considered as salts of weak hydrogen sulfide acid (aqueous solution of H2S). Many S. are natural ... ... Big Polytechnic Encyclopedia
Oh; mn (units sulphide, and; m.). [from lat. sulfur sulfur] Chem. Sulfur compounds with metals and some non-metals. C. titanium. Organic with Natural s. (class of minerals). ◁ Sulfide, th, oh. With mixed mixtures. From ore ore. * * * sulfides of sulfur compounds ... encyclopedic Dictionary
- (from the Latin sulfur sulfur) a class of chemical compounds that are metal compounds (as well as a number of non-metals B, Si, P, As) with sulfur (S), where it has an oxidation state of −2. Can be considered as salts of hydrogen sulfide ... ... Wikipedia
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- Uranium deposits of the type of disagreement in the Phanerozoic, Andrei Borisovich Khalezov. The features of uranium deposits in the structures of river paleo-valley on the surfaces of structural and stratigraphic disagreement are considered on the example of three regions of the Ural region with different…
Sulfur - The element of the 3rd period and the VIA group of the Periodic system, serial number 16, refers to chalcogens. The electronic formula of the atom [10 Ne] 3s 2 3p 4, characteristic oxidation states 0, ‑II, + IV and + VI, the state S VI is considered stable.
Sulfur oxidation scale:
The electronegativity of sulfur is 2.60, it is characterized by non-metallic properties. In hydrogen and oxygen compounds it is contained in various anions, it forms oxygen-containing acids and their salts, binary compounds.
In nature - fifteenth chemical prevalence element (the seventh among non-metals). It is found in the free (native) and related form. A vital element for higher organisms.
Sulfur S. Simple substance. Crystalline yellow (α-rhombic and β-monoclinic,
at 95.5 ° C) or amorphous (plastic). In the lattice sites there are S 8 molecules (non-planar “corona” type cycles), amorphous sulfur consists of S n chains. Low-melting substance, the viscosity of the liquid passes through a maximum at 200 ° C (the gap of molecules S 8, the interlacing of chains S n). In a pair - molecules S 8, S 6, S 4, S 2. At 1500 ° C, monoatomic sulfur appears (in chemical equations for simplicity, any sulfur is represented as S).
Sulfur does not dissolve in water and under normal conditions does not react with it, it is well soluble in carbon disulfide CS 2.
Sulfur, especially powdered, has a high activity when heated. Reacts as an oxidizing agent with metals and non-metals:
but as reducing agent - with fluorine, oxygen and acids (when 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 + 2НI (g) = I 2 + H 2 S,
but in solution the reaction goes the opposite way:
I 2 + H 2 S (p) = 2 HI + S ↓
Getting: at industry smelted from natural deposits of native sulfur (using water vapor), released during desulfurization of coal gasification products.
Sulfur is used to synthesize carbon disulfide, sulfuric acid, sulphurous (vat) dyes, in the vulcanization of rubber, as a means of protecting plants from powdery mildew, and for treating skin diseases.
Hydrogen sulfide H 2 S. Oxygen free acid. Colorless gas with a suffocating smell, heavier than air. The molecule has the structure of a double-incomplete tetrahedron [:: S (H) 2]
(sp 3 ‑ hybridization, the H – S – H jack angle is far from tetrahedral). Unstable when heated above 400 ° C. It is slightly soluble in water (2.6 l / 1 l H 2 O at 20 ° C), a saturated solution is decimolar (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 -1 is equal to 1 10 –13 mol / l). When standing in air, the solution becomes cloudy (inhibitor - sucrose). It is neutralized by alkalis, not completely - by ammonia hydrate. Strong reducing agent. It enters into ion exchange reactions. Sulfiding agent, precipitates from a solution of differently colored sulphides with very low solubility.
Qualitative reactions - sedimentation of sulphides, as well as incomplete combustion of H 2 S with the formation of a yellow sulfur deposit on a cold object introduced into the flame (porcelain spatula). A byproduct of refining petroleum, natural and coke oven gas.
It is used in the production of sulfur, inorganic and organic sulfur-containing compounds as an analytical reagent. Extremely poisonous. The equations of the most important reactions:
Getting: at industry - direct synthesis:
H 2 + S = H 2 S(150–200 ° C)
or by heating sulfur with paraffin;
at laboratories - displacement of sulfides by strong acids
FeS + 2НCl (conc.) = FeCl 2 + H 2 S
or complete hydrolysis of binary compounds:
Al 2 S 3 + 6H 2 O = 2Al (OH) 3 ↓ + 3 H 2 S
Sodium sulfide Na 2 S. Oxygen free salt. White, very hygroscopic. Melts without decomposition, thermally stable. It is well soluble in water, hydrolyzed by anion, creates a highly alkaline medium in solution. When standing in air, the solution becomes cloudy (colloidal sulfur) and turns yellow (polysulfide dye). Typical reducing agent. Attaches sulfur. It enters into ion exchange reactions.
Qualitative reactions on ion S 2‑– deposition of variously colored metal sulfides, of which MnS, FeS, ZnS are decomposed into HCl (s).
It is used in the production of sulfur dyes and cellulose, to remove the hair of hides during tanning of leather, as a reagent in analytical chemistry.
The equations of the most important reactions:
Na 2 S + 2НCl (dec.) = 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 (sat.) = 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 (corporal) ↓; FeS (black) ↓; ZnS (white) ↓
S 2‑ + 2Ag + = Ag 2 S (black) ↓
S 2‑ + M 2+ = CdS (yellow) ↓; PbS, CuS, HgS (black) ↓
3S 2‑ + 2Bi 3+ = Bi 2 S 3 (corr. - black) ↓
3S 2‑ + 6H 2 O + 2M 3+ = 3H 2 S + 2M (OH) 3 ↓ (M = Al, Cr)
Getting at 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 + 4C (coke) = Na 2 S + 4CO (800–1000 ° C)
Na 2 SO 4 + 4CO = Na 2 S + 4СO 2 (600–700 ° C)
Aluminum sulfide Al 2 S 3. Oxygen free salt. White, the Al - S bond is predominantly covalent. It melts without decomposition under the excessive pressure of N 2, it is easily sublimated. It is oxidized in air when calcined. Fully hydrolyzed with water, does not precipitate from solution. Decomposed by strong acids. It is used as a solid source of pure hydrogen sulfide. The equations of the most important reactions:
Al 2 S 3 + 6H 2 O = 2Al (OH) 3 ↓ + 3H 2 S (pure)
Al 2 S 3 + 6НCl (dec.) = 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)
Getting: 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 and gray with a green tint, refractory, decomposes when heated in a vacuum. In the wet state is sensitive to oxygen in the air. Insoluble in water. Does not precipitate upon saturation of solutions of iron (II) salts with hydrogen sulfide. Decomposed by acids. It is used as a raw material in the production of iron, a solid source of hydrogen sulfide.
The compound of iron (III) composition of Fe 2 S 3 not known (not received).
The equations of the most important reactions:
Receiving:
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 on ignition. Insoluble in water, does not react with diluted acids, alkalis. It is decomposed by acid-oxidizing agents, roasted in air. It is used as a raw material in the production of iron, sulfur and sulfuric acid, a catalyst in organic synthesis. In nature - ore minerals pyrites and marcasite.
The equations of the most important reactions:
FeS 2 = FeS + S (above 1170 ° C, vacuum)
2FeS 2 + 14H 2 SO 4 (conc., Mountains) = Fe 2 (SO 4) 3 + 15SO 2 + 14Н 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. Oxygen free acid salt. White, melted under pressure. Extremely volatile, thermally unstable. It oxidizes on air. It is well soluble in water, hydrolyzed by cation and anion (prevails), creates an alkaline environment. The solution turns yellow in the air. It is decomposed by acids, in a saturated solution it adds sulfur. Alkalis are not neutralized, the middle salt (NH 4) 2 S does not exist in the solution (for conditions of obtaining the middle salt, see the heading “H 2 S”). It is used as a component of photo developers, as an analytical reagent (sulphide precipitator).
The equations of the most important reactions:
NH 4 HS = NH 3 + H 2 S (above 20 ° C)
NH 4 HS + HCl (dec.) = 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 (sat. H 2 S) + 2CuSO 4 = (NH 4) 2 SO 4 + H 2 SO 4 + 2CuS ↓
Getting: saturation of 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 as a solution of (NH 4) 2 S and the average salt formula is used in the writing of the reaction equations, although ammonium sulfide is completely hydrolyzed in water to NH 4 HS and NH 3 H 2 O.
Sulphur dioxide. Sulfites
Sulfur dioxide SO 2. Acid 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. It is soluble in water (~ 40 l / 1 l H 2 O at 20 ° C). Forms a polyhydrate with the properties of a weak acid, the products of dissociation - ions HSO 3 - and SO 3 2−. Ion HSO 3 - has two tautomeric forms - symmetrical (non-acidic) with the structure of the tetrahedron (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 ion SO 3 2‑ is also tetrahedral [: S (O) 3].
Reacts with alkalis, ammonia hydrate. Typical reducing agent, weak oxidizing agent.
Qualitative reaction - discoloration of yellow-brown "iodine water". Intermediate in the production of sulfites and sulfuric acid.
It is used for bleaching wool, silk and straw, preserving and storing fruits, as a disinfectant, antioxidant, coolant. Poisonous.
The compound of H 2 SO 3 (sulfurous acid) is not known (does not exist).
The equations of the most important reactions:
Dissolution in water and acidic properties:
Getting: in industry - the burning of sulfur in air enriched with oxygen, and, to a lesser extent, the roasting of sulphide ores (SO 2 - associated gas during roasting of pyrite):
S + O 2 = SO 2(280–360 ° C)
4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8 SO 2(800 ° C, roasting)
in the laboratory - the 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. Saline White. When heated in air it decomposes without melting, melts under argon overpressure. In the wet state and in solution sensitive to oxygen in the air. It is soluble in water, hydrolyzed by anion. Decomposed by acids. Typical reducing agent.
Qualitative reaction on ion SO 3 2‑ - the formation of a white precipitate of barium sulfite, which is transferred into the solution with strong acids (HCl, HNO 3).
It is used as a reagent in analytical chemistry, a component of photographic solutions, a neutralizer of chlorine in the bleaching of fabrics.
The equations of the most important reactions:
Receiving:
Na 2 CO 3 (conc.) + SO 2 = Na 2 SO 3 + CO 2
Sulphuric acid. Sulfates
Sulfuric acid H 2 SO 4. Oxoic acid. Colorless liquid, very viscous (oily), very hygroscopic. The molecule has a distorted тет tetrahedral structure (sp 3 гиб hybridization), contains covalent σ ‑ bonds S - OH and σπ ‑ bonds S = O. The ion SO 4 2‑ has a correctly ‑ tetrahedral structure. It has a wide temperature range of the liquid state (~ 300 degrees). When heated above 296 ° C it partially decomposes. It is distilled as an azeotropic mixture with water (mass fraction of acid is 98.3%, boiling point is 296–340 ° C), decomposes completely with stronger heating. Unlimited mixes with water (with strong exo-Effect). Strong acid in solution, neutralized by alkali and ammonia hydrate. Converts metals to sulphates (with an excess of concentrated acid under normal conditions, soluble hydrosulphates are formed), but Be, Bi, Co, Fe, Mg and Nb metals are passivated in the concentrated acid and do not react with it. Reacts with basic oxides and hydroxides, decomposes salts of weak acids. Weak oxidizer in a dilute solution (due to H I), strong - in a concentrated solution (due to S VI). Well dissolves SO 3 and reacts with it (a heavy oily liquid is formed - oleum, contains H 2 S 2 O 7).
Qualitative reaction on SO 4 2‑ ion - precipitation of white barium sulphate BaSO 4 (the precipitate is not converted into a solution of hydrochloric and nitric acids, in contrast to the white precipitate BaSO 3).
It is used in the production of sulphates and other sulfur compounds, mineral fertilizers, explosives, dyes and drugs, in organic synthesis, for “opening” (the first stage of processing) of industrially important ores and minerals, in refining petroleum products, electrolysis of water, as electrolyte of lead batteries . Poisonous, causes skin burns. The equations of the most important reactions:
Getting at 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, roasting)
2H 2 S + 3O 2 (g) = 2 SO 2 + 2H 2 O (250–300 ° C)
CaSO 4 + С (coke) = CaO + SO 2 + CO (1300–1500 ° C)
b) the conversion of SO 2 to SO 3 in the contact apparatus:
c) synthesis of concentrated and anhydrous sulfuric acid:
H 2 O (Coll. H 2 SO 4) + SO 3 = H 2 SO 4(conc., anhyd.)
(the absorption of SO 3 by pure water with the production of 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, disaric acid H 2 S 2 O 7 and excess SO 3. Dissolved SO 3 ensures oleum waterlessness (H 2 SO 4 is immediately formed when water enters), which allows it to be safely transported in steel tanks.
Sodium sulfate Na 2 SO 4. Saline White, hygroscopic. Melts and boils without decomposition. Forms crystalline hydrate (mineral mirabilite) easily losing water; technical name Glauber's salt. It is soluble in water, not hydrolyzed. Reacts with H 2 SO 4 (conc.), SO 3. Restored by hydrogen, coke when heated. It enters into ion exchange reactions.
It is used in the manufacture of glass, cellulose and mineral paints, as a drug. Contained in brine of salt lakes, in particular in the Gulf of Kara-Bogaz-Gol of the Caspian Sea.
The equations of the most important reactions:
Potassium hydrosulfate KHSO 4. Sour oxol salt. White, hygroscopic, but does not form crystalline hydrates. When heated melts and decomposes. It is well soluble in water, the anion undergoes dissociation in solution, the solution medium is strongly acid. It is neutralized by alkalis.
It is used as a component of fluxes in metallurgy, an integral part of mineral fertilizers.
The 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)
Getting: treatment of potassium sulfate in the cold (with more than 6O%) sulfuric acid
K 2 SO 4 + H 2 SO 4 (conc.) = 2 KHSO 4
Calcium sulfate CaSO 4. Saline White, very hygroscopic, refractory, decomposes when calcined. Natural CaSO 4 is found in the form of a very common mineral gypsum CaSO 4 2H 2 O. At 130 ° C, gypsum loses some of the water and goes into burnt plaster2CaSO 4 H 2 O (technical name alabaster). Fully dehydrated (200 ° C) gypsum responds to the mineral anhydriteCaSO 4. It is slightly soluble in water (0.206 g / 100 g H 2 O at 20 ° C), the solubility decreases when heated. Reacts with H 2 SO 4 (conc.). Recovered by coke during fusion. Defines most of the "constant" hardness of fresh water (for more details, see 9.2).
The 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, paper filler. The binding mortar prepared from burnt gypsum “sets” faster than the mixture based on Ca (OH) 2. Hardening is provided by the binding of water, the formation of gypsum in the form of stone mass. Burnt gypsum is used to make plaster casts, architectural and decorative forms and products, partition walls and panels, and stone floors.
Aluminum potassium sulfate KAl (SO 4) 2. Double oxol salt. White, hygroscopic. With strong heat decomposes. Forms crystalline hydrate - alumina potassium alum. Moderately soluble in water, hydrolyzed by aluminum cation. Reacts with alkalis, ammonia hydrate.
It is used as a mordant in dyeing fabrics, tanning leather, coagulant in fresh water purification, a component of paper sizing compositions, and an external hemostatic agent in medicine and cosmetology. It is formed by the co-crystallization of aluminum and potassium sulfates.
The equations of the most important reactions:
Chromium (III) sulfate - potassium KCr (SO 4) 2. Double oxol salt. Red (dark purple hydrate, technical name potassium alum). When heated decomposes without melting. It is readily soluble in water (the gray-blue color of the solution corresponds to the aquacomplex 3+), it is hydrolyzed by the chromium (III) cation. Reacts with alkalis, ammonia hydrate. Weak oxidizer and reducing agent. It enters into ion exchange reactions.
Qualitative reactions to Cr 3+ ion - reduction to Cr 2+ or oxidation to yellow CrO 4 2‑.
It is used as a tanning agent for leather, a mordant in dyeing fabrics, a reagent in a photograph. It is formed by the co-crystallization of chromium (III) and potassium sulphates. The equations of the most important reactions:
Manganese (II) sulfate MnSO 4. Saline White, melts and decomposes on ignition. Crystalline hydrate MnSO 4 5H 2 O - red ‑ pink, technical name manganese vitriol. It is soluble in water, light pink (almost colorless) color of the solution corresponds to the aquacomplex 2+; hydrolyzed by cation. Reacts with alkalis, ammonia hydrate. Weak reducing agent reacts with typical (strong) oxidizing agents.
Qualitative reactions to the Mn 2+ ion — a conduction with the MnO 4 ion and the disappearance of the violet color of the latter, the oxidation of Mn 2+ to MnO 4, and the appearance of a violet color.
It is used to obtain Mn, MnO 2 and other compounds of manganese, as micronutrient and analytical reagent.
The equations of the most important reactions:
Receiving:
2MnO 2 + 2H 2 SO 4 (conc.) = 2 MnSO 4 + O 2 + 2H 2 O (100 ° C)
Iron (II) sulfate FeSO 4. Saline White (hydrate light green, technical name inkstone),hygroscopic. It decomposes when heated. It is soluble in water, to a small extent hydrolyzed by cation. It is rapidly oxidized in solution by oxygen in the air (the solution turns yellow and becomes cloudy). Reacts with acid-oxidizing agents, alkalis, ammonia hydrate. Typical reducing agent.
It is used as a component of mineral paints, electrolytes in electroplating, wood preservative, fungicide, anti-anemia drug. In the laboratory, more often it is taken as the double salt Fe (NH 4) 2 (SO 4) 2 6Н 2 O ( salt mora) more resistant to air.
The equations of the most important reactions:
Receiving:
Fe + H 2 SO 4 (dec.) = FeSO 4+ H 2
FeCO 3 + H 2 SO 4 (par.) = FeSO 4 + CO 2 + H 2 O
7.4. VA-group non-metals
Nitrogen. Ammonia
Nitrogen - an element of the 2nd period and the VA ‑ group of the Periodic system, serial number 7. Atomic electron formula [2 He] 2s 2 2p 3, characteristic oxidation states 0, –III, + III and + V, less often + II, + IV and others; N v state is considered relatively stable.
Nitrogen Oxidation Scale:
Nitrogen has a high electronegativity (3.07), the third after F and O. Shows 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 chemical prevalence 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, odorless gas that condenses into a colorless liquid (unlike O 2).
The main component of the air: 78.09% by volume, 75.52% by weight. From liquid air, nitrogen boils over oxygen O 2. It is 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, N 2 reacts only with lithium (in a humid atmosphere), forming lithium nitride Li 3 N, nitrides of other elements are synthesized with strong heating:
N 2 + 3Mg = Mg 3 N 2 (800 ° C)
In an electrical discharge, N 2 reacts with fluorine and to a very small extent with oxygen:
A reversible reaction to produce ammonia takes place at 500 ° C, under pressure up to 350 atm and necessarily in the presence of a catalyst (Fe / F 2 O 3 / FeO, in the Pt laboratory):
In accordance with the Le Chatelier principle, an increase in the ammonia yield should occur with an increase in pressure and a decrease in temperature. However, the reaction rate at low temperatures is very low, so the process is carried out at 450–500 ° C, reaching a 15% yield of ammonia. Unreacted N 2 and H 2 return to the reactor and thereby increase the degree of reaction.
Nitrogen is chemically passive with respect to acids and alkalis, does not support combustion.
Getting at industry - fractional distillation of liquid air or the removal of oxygen from air by chemical means, for example, by the reaction 2C (coke) + O 2 = 2CO by heating. In these cases, nitrogen is produced, which also contains impurities of noble gases (mainly argon).
AT laboratories small amounts of chemically pure nitrogen can be obtained by a combination 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 environment for chemical and metallurgical processes and storage of flammable substances.
Ammonia NH 3. Binary compound, the degree of oxidation 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 гиб-hybridization). The presence of nitrogen in the NH 3 molecule of an electron donor pair in a sp 3 ‑ hybrid orbital causes a characteristic addition reaction of the hydrogen cation, with the formation of a cation ammonium NH 4 +. Liquefied under pressure at room temperature. In the liquid state is associated due to hydrogen bonds. Thermally unstable. It is soluble in water (more than 700 l / 1 l of H 2 O at 20 ° C); the proportion in the saturated solution is = 34% by mass and = 99% by volume, pH = 11.8.
Highly reactive, prone to addition reactions. Arises in oxygen, reacts with acids. It exhibits reducing (due to N ‑III) and oxidizing (due to H I) properties. Dried out only with calcium oxide.
Qualitative reactions - the formation of white "smoke" in contact with gaseous HCl, blackening of the paper moistened with a solution of Hg 2 (NO 3) 2.
Intermediate product in the synthesis of HNO 3 and ammonium salts. Used in the manufacture of soda, nitrogen fertilizers, dyes, explosives; liquid ammonia is a coolant. Poisonous.
The equations of the most important reactions:
Getting: at laboratories - displacement of ammonia from ammonium salts when heated with soda lime (NaOH + CaO):
or boiling an aqueous solution of ammonia, followed by drying the gas.
AT industry ammonia is synthesized from nitrogen (see) with hydrogen. Produced by the industry either in a 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 compound. White, in the crystal lattice are NH 3 and H 2 O molecules, bound by a weak hydrogen bond H 3 N ... HON. Present in an aqueous solution of ammonia, a weak base (the products of dissociation - the cation NH 4 - and the anion OH -). The ammonium cation has a regularly π-tetrahedral structure (sp 3 ‑ hybridization). Thermally unstable, completely decomposed by boiling the solution. Neutralized by strong acids. Shows reducing properties (due to N III) in a concentrated solution. It enters into the reaction of ion exchange and complexation.
Qualitative reaction - the formation of white "smoke" in contact with gaseous HCl.
It is used to create a weak alkaline medium in solution, during the precipitation of amphoteric hydroxides.
The 1M ammonia solution contains mainly NH 3 H 2 O hydrate and only 0.4% of NH 4 + and OH - ions (due to the dissociation of the hydrate); thus, the ionic “ammonium hydroxide NH 4 OH” is practically not contained in the solution, there is no such compound in the solid hydrate. The 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 (dec.) = 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.) + ЗBr 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 + 2KON
4 (NH 3 H 2 O) (conc.) + Ag 2 O = 2OH + 3H 2 O
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
Diluted ammonia solution (3-10% ‑ ‑) is often called liquid ammonia (the name was invented by the alchemists), and the concentrated solution (18.5–25%) ammonia water (produced by industry).