Sulfur oxide (sulfur dioxide, sulfur dioxide, sulfur dioxide) is a colorless gas that has normal conditions a sharp characteristic odor (similar to the smell of a burning match). Liquefied under pressure room temperature. Sulfur dioxide is soluble in water, and unstable sulfuric acid is formed. This substance is also soluble in sulfuric acid and ethanol. This is one of the main components that make up volcanic gases.
Sulphur dioxide
The production of SO2 - sulfur dioxide - industrially involves burning sulfur or roasting sulfides (mainly pyrite is used).
4FeS2 (pyrite) + 11O2 = 2Fe2O3 + 8SO2 (sulfur dioxide).
In a laboratory setting, sulfur dioxide can be produced by treating hydrosulfites and sulfites with strong acids. In this case, the resulting sulfurous acid immediately breaks down into water and sulfur dioxide. For example:
Na2SO3 + H2SO4 (sulfuric acid) = Na2SO4 + H2SO3 (sulfurous acid).
H2SO3 (sulfurous acid) = H2O (water) + SO2 (sulfur dioxide).
The third method of producing sulfur dioxide involves the action of concentrated sulfuric acid on low-active metals when heated. For example: Cu (copper) + 2H2SO4 (sulfuric acid) = CuSO4 (copper sulfate) + SO2 (sulfur dioxide) + 2H2O (water).
Chemical properties sulfur dioxide
Formula sulfur dioxide- SO3. This substance belongs to acid oxides.
1. Sulfur dioxide dissolves in water, resulting in sulfurous acid. IN normal conditions this reaction is reversible.
SO2 (sulfur dioxide) + H2O (water) = H2SO3 (sulfurous acid).
2. With alkalis, sulfur dioxide forms sulfites. For example: 2NaOH (sodium hydroxide) + SO2 (sulfur dioxide) = Na2SO3 (sodium sulfite) + H2O (water).
3. The chemical activity of sulfur dioxide is quite high. The reducing properties of sulfur dioxide are most pronounced. In such reactions, the oxidation state of sulfur increases. For example: 1) SO2 (sulfur dioxide) + Br2 (bromine) + 2H2O (water) = H2SO4 (sulfuric acid) + 2HBr (hydrogen bromide); 2) 2SO2 (sulfur dioxide) + O2 (oxygen) = 2SO3 (sulfite); 3) 5SO2 (sulfur dioxide) + 2KMnO4 (potassium permanganate) + 2H2O (water) = 2H2SO4 (sulfuric acid) + 2MnSO4 (manganese sulfate) + K2SO4 (potassium sulfate).
The last reaction is an example of a qualitative reaction to SO2 and SO3. The solution becomes purple in color.)
4. In the presence of strong reducing agents, sulfur dioxide can exhibit oxidizing properties. For example, in order to extract sulfur from exhaust gases in the metallurgical industry, they use the reduction of sulfur dioxide with carbon monoxide (CO): SO2 (sulfur dioxide) + 2CO (carbon monoxide) = 2CO2 + S (sulfur).
Also, the oxidizing properties of this substance are used to obtain phosphorous acid: PH3 (phosphine) + SO2 (sulfur dioxide) = H3PO2 (phosphoric acid) + S (sulfur).
Where is sulfur dioxide used?
Sulfur dioxide is mainly used to produce sulfuric acid. It is also used in the production of low-alcohol drinks (wine and other medium-sized drinks price category). Due to the property of this gas to kill various microorganisms, it is used to fumigate warehouses and vegetable stores. In addition, sulfur oxide is used to bleach wool, silk, and straw (those materials that cannot be bleached with chlorine). In laboratories, sulfur dioxide is used as a solvent and in order to obtain various salts of sulfur dioxide.
Physiological effects
Sulfur dioxide has strong toxic properties. Symptoms of poisoning are cough, runny nose, hoarseness, a peculiar taste in the mouth, and severe sore throat. When sulfur dioxide is inhaled in high concentrations, difficulty swallowing and choking, speech disturbance, nausea and vomiting occur, and acute pulmonary edema may develop.
MPC of sulfur dioxide:
- indoors - 10 mg/m³;
- average daily maximum one-time per atmospheric air- 0.05 mg/m³.
Sensitivity to sulfur dioxide varies among individuals, plants, and animals. For example, among trees the most resistant are oak and birch, and the least resistant are spruce and pine.
Structure of the SO2 molecule
The structure of the SO2 molecule is similar to the structure of the ozone molecule. The sulfur atom is in a state of sp2 hybridization, the shape of the orbitals is a regular triangle, and the shape of the molecule is angular. The sulfur atom has a lone pair of electrons. The S–O bond length is 0.143 nm, and the bond angle is 119.5°.
The structure corresponds to the following resonant structures:
Unlike ozone, the multiplicity of the S – O bond is 2, that is, the main contribution is made by the first resonant structure. The molecule is characterized by high thermal stability.
Sulfur compounds +4 - exhibit redox duality, but with a predominance of reducing properties.
1. Interaction of SO2 with oxygen
2S+4O2 + O 2 S+6O
2. When SO2 is passed through hydrogen sulfide acid, sulfur is formed.
S+4O2 + 2H2S-2 → 3So + 2 H2O
4 S+4 + 4 → So 1 - oxidizing agent (reduction)
S-2 - 2 → So 2 - reducing agent (oxidation)
3. Sulfurous acid is slowly oxidized by atmospheric oxygen into sulfuric acid.
2H2S+4O3 + 2O → 2H2S+6O
4 S+4 - 2 → S+6 2 - reducing agent (oxidation)
O + 4 → 2O-2 1 - oxidizing agent (reduction)
Receipt:
1) sulfur (IV) oxide in industry:
sulfur combustion:
pyrite firing:
4FeS2 + 11O2 = 2Fe2O3
in the laboratory:
Na2SO3 + H2SO4 = Na2SO4 + SO2 + H2O
Sulphur dioxide, preventing fermentation, facilitates the deposition of pollutants, scraps of grape tissue with pathogenic microflora and allows alcoholic fermentation to be carried out using pure yeast cultures in order to increase the yield of ethyl alcohol and improve the composition of other alcoholic fermentation products.
The role of sulfur dioxide is thus not limited to antiseptic actions that improve the environment, but also extends to improving technological conditions fermentation and storage of wine.
These conditions are correct use sulfur dioxide (limiting the dosage and time of contact with air) lead to an increase in the quality of wines and juices, their aroma, taste, as well as transparency and color - properties associated with the resistance of wine and juice to turbidity.
Sulfur dioxide is the most common air pollutant. It is released by all power plants when burning fossil fuels. Sulfur dioxide can also be emitted by metallurgical industry enterprises (source: coking coals), as well as nearby chemical production(for example, production of sulfuric acid). It is formed during the decomposition of sulfur-containing amino acids that were part of the proteins of ancient plants that formed deposits of coal, oil, and oil shale.
Finds application in industry for bleaching various products: cloth, silk, paper pulp, feathers, straw, wax, bristles, horsehair, food products, for the disinfection of fruits and canned food, etc. As a by-product, sulfur dioxide is formed and released into the air of work areas in a number of industries: sulfuric acid, cellulose, during roasting of ores containing sulfur metals, in pickling rooms at metal plants In the production of glass, ultramarine, etc., sulfur is often found in the air of boiler rooms and ash rooms, where it is formed during the combustion of sulfur-containing coals.
When dissolved in water, a weak and unstable sulfurous acid H2SO3 (exists only in aqueous solution)
SO2 + H2O ↔ H2SO3
Sulfurous acid dissociates stepwise:
H2SO3 ↔ H+ + HSO3- (first step, hydrosulfite anion is formed)
HSO3- ↔ H+ + SO32- (second stage, sulfite anion is formed)
H2SO3 forms two series of salts - medium (sulfites) and acidic (hydrosulfites).
A qualitative reaction to salts of sulfurous acid is the interaction of the salt with a strong acid, which releases SO2 gas with a pungent odor:
Na2SO3 + 2HCl → 2NaCl + SO2 + H2O 2H+ + SO32- → SO2 + H2O
In this article you will find information about what sulfur oxide is. Its basic chemical and physical properties, existing forms, methods of their preparation and differences from each other will be considered. The applications and biological role of this oxide in its various forms will also be mentioned.
What is the substance
Sulfur oxide is a compound of simple substances, sulfur and oxygen. There are three forms of sulfur oxides, differing in the degree of valence S, namely: SO (sulfur monoxide, sulfur monoxide), SO 2 (sulfur dioxide or sulfur dioxide) and SO 3 (sulfur trioxide or anhydride). All of the listed variations of sulfur oxides have similar chemical and physical characteristics.
General information about sulfur monoxide
Divalent sulfur monoxide, or otherwise sulfur monoxide, is an inorganic substance consisting of two simple elements- sulfur and oxygen. Formula - SO. Under normal conditions, it is a colorless gas, but with a pungent and specific odor. Reacts with an aqueous solution. Quite a rare connection in earth's atmosphere. It is unstable to temperature and exists in dimeric form - S 2 O 2 . Sometimes it is capable of interacting with oxygen to form sulfur dioxide as a result of the reaction. Does not form salts.
Sulfur oxide (2) is usually obtained by burning sulfur or decomposing its anhydride:
- 2S2+O2 = 2SO;
- 2SO2 = 2SO+O2.
The substance dissolves in water. As a result, sulfur oxide forms thiosulfuric acid:
- S 2 O 2 + H 2 O = H 2 S 2 O 3 .
General data on sulfur dioxide
Sulfur oxide is another form of sulfur oxides with chemical formula SO2. It has an unpleasant specific odor and is colorless. When subjected to pressure, it can ignite at room temperature. When dissolved in water, it forms unstable sulfurous acid. Can dissolve in ethanol and sulfuric acid solutions. It is a component of volcanic gas.
In industry it is obtained by burning sulfur or roasting its sulfides:
- 2FeS 2 +5O 2 = 2FeO+4SO 2.
In laboratories, as a rule, SO 2 is obtained using sulfites and hydrosulfites, exposing them to strong acid, as well as to exposure of metals with a low degree of activity to concentrated H 2 SO 4.
Like other sulfur oxides, SO2 is an acidic oxide. Interacting with alkalis, forming various sulfites, it reacts with water, creating sulfuric acid.
SO 2 is extremely active, and this is clearly expressed in its reducing properties, where the oxidation state of sulfur oxide increases. May exhibit oxidizing properties if exposed to a strong reducing agent. The last one characteristic feature used for the production of hypophosphorous acid, or for the separation of S from gases in the metallurgical field.
Sulfur oxide (4) is widely used by humans to produce sulfurous acid or its salts - this is its main area of application. It also participates in winemaking processes and acts there as a preservative (E220); sometimes it is used to pickle vegetable stores and warehouses, as it destroys microorganisms. Materials that cannot be bleached with chlorine are treated with sulfur oxide.
SO 2 is a rather toxic compound. Characteristic symptoms symptoms indicating poisoning by it are coughing, breathing problems, usually in the form of a runny nose, hoarseness, the appearance of an unusual taste and a sore throat. Inhalation of such gas can cause suffocation, impaired speech ability of the individual, vomiting, difficulty swallowing, and acute pulmonary edema. The maximum permissible concentration of this substance in the work area is 10 mg/m3. However, different people's bodies may exhibit different sensitivity to sulfur dioxide.
General information about sulfuric anhydride
Sulfur gas, or sulfuric anhydride as it is called, is a higher oxide of sulfur with the chemical formula SO 3 . Liquid with a suffocating odor, highly volatile under standard conditions. It is capable of solidifying, forming crystalline mixtures from its solid modifications, at temperatures of 16.9 °C and below.
Detailed analysis of higher oxide
When SO 2 is oxidized by air under the influence of high temperatures, a necessary condition is the presence of a catalyst, for example V 2 O 5, Fe 2 O 3, NaVO 3 or Pt.
Thermal decomposition of sulfates or interaction of ozone and SO 2:
- Fe 2 (SO 4)3 = Fe 2 O 3 +3SO 3;
- SO 2 +O 3 = SO 3 +O 2.
Oxidation of SO 2 with NO 2:
- SO 2 +NO 2 = SO 3 +NO.
To physical quality characteristics include: the presence in the gas state of a flat structure, trigonal type and D 3 h symmetry; during the transition from gas to crystal or liquid, it forms a trimer of a cyclic nature and a zigzag chain, has a covalent polar bond.
In solid form, SO 3 occurs in alpha, beta, gamma and sigma forms, and it has, accordingly, different melting points, degrees of polymerization and a variety of crystalline forms. The existence of such a number of SO 3 species is due to the formation of donor-acceptor type bonds.
The properties of sulfur anhydride include many of its qualities, the main ones being:
Ability to interact with bases and oxides:
- 2KHO+SO 3 = K 2 SO 4 +H 2 O;
- CaO+SO 3 = CaSO 4.
Higher sulfur oxide SO3 has quite a high activity and creates sulfuric acid by interacting with water:
- SO 3 + H 2 O = H2SO 4.
It reacts with hydrogen chloride and forms chlorosulfate acid:
- SO 3 +HCl = HSO 3 Cl.
Sulfur oxide is characterized by the manifestation of strong oxidizing properties.
Sulfuric anhydride is used in the creation of sulfuric acid. A small amount of it is released into environment while using sulfur bombs. SO 3, forming sulfuric acid after interaction with a wet surface, destroys a variety of dangerous organisms, such as fungi.
Summing up
Sulfur oxide can be in different states of aggregation, ranging from liquid to solid form. It is rare in nature, but there are quite a few ways to obtain it in industry, as well as areas where it can be used. The oxide itself has three forms in which it exhibits different degrees of valence. Can be very toxic and cause serious problems with health.
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Sulfur. Hydrogen sulfide, sulfides, hydrosulfides. Sulfur oxides (IV) and (VI). Sulfurous and sulfuric acids and their salts. Esters of sulfuric acid. Sodium thiosulfate
4.1. Sulfur
Sulfur is one of the few chemical elements that people have been using for several millennia. It is widespread in nature and is found both in a free state (native sulfur) and in compounds. Minerals containing sulfur can be divided into two groups - sulfides (pyrites, lusters, blende) and sulfates. Native sulfur is found in large quantities in Italy (the island of Sicily) and the USA. In the CIS there are deposits of native sulfur in the Volga region, in the states Central Asia, in Crimea and other areas.
Minerals of the first group include lead luster PbS, copper luster Cu 2 S, silver luster - Ag 2 S, zinc blende - ZnS, cadmium blende - CdS, pyrite or iron pyrite - FeS 2, chalcopyrite - CuFeS 2, cinnabar - HgS.
Minerals of the second group include gypsum CaSO 4 2H 2 O, mirabilite (Glauber's salt) - Na 2 SO 4 10H 2 O, kieserite - MgSO 4 H 2 O.
Sulfur is found in the bodies of animals and plants, as it is part of protein molecules. Organic sulfur compounds are found in oil.
Receipt
1. When obtaining sulfur from natural compounds, for example from sulfur pyrites, 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 with a lack of oxygen according to the reaction:
2H 2 S+O 2 =2S+2H 2 O
3. Currently, it is common to obtain sulfur by reducing sulfur dioxide SO2 with carbon - a by-product in the smelting of metals from sulfur ores:
SO 2 +C = CO 2 +S
4. Exhaust gases from metallurgical and coke ovens contain a mixture of sulfur dioxide and hydrogen sulfide. This mixture is passed through high temperature above the catalyst:
H 2 S+SO 2 =2H 2 O+3S
^ Physical properties
Sulfur is a hard, brittle, lemon-yellow substance. It is practically insoluble in water, but is highly soluble in carbon disulfide CS 2 aniline and some other solvents.
Conducts heat poorly and electricity. Sulfur forms several allotropic modifications:
1 . ^ Rhombic sulfur (the most stable), the 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 rises, it darkens and becomes so viscous that it does not flow out of the test tube; with further heating, the viscosity drops again, and at 444.6 °C sulfur boils.
2. ^ Monoclinic sulfur - modification in the form of dark yellow needle-shaped crystals, obtained by slow cooling of molten sulfur.
3. Plastic sulfur is formed if sulfur heated to a boil is poured into cold water. Easily stretches 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
The sulfur atom, having an incomplete external energy level, can add two electrons and exhibit a degree
Oxidation -2. Sulfur exhibits this degree of oxidation in compounds with metals and hydrogen (Na 2 S, H 2 S). When electrons are given away or withdrawn to an atom of a more electronegative element, the oxidation state of sulfur can be +2, +4, +6.
In the cold, sulfur is relatively inert, but with increasing temperature its reactivity increases. 1. With metals, sulfur exhibits oxidizing properties. These reactions produce sulfides (does not react with gold, platinum and iridium): Fe+S=FeS
2. Under normal conditions, sulfur does not interact with hydrogen, but at 150-200°C a reversible reaction occurs:
3. In reactions with metals and hydrogen, sulfur behaves as a typical oxidizing agent, and in the presence of strong oxidizing agents it exhibits reducing properties.
S+3F 2 =SF 6 (does not react with iodine)
4. Combustion of sulfur in oxygen occurs at 280°C, and in air at 360°C. This produces a mixture of SO 2 and SO 3:
S+O 2 =SO 2 2S+3O 2 =2SO 3
5. When heated without air access, sulfur directly combines with phosphorus and carbon, exhibiting 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. Thus, when sulfur powder is boiled with alkalis, sulfites and sulfides are formed:
Application
Sulfur is widely used in industry and agriculture. About half of its production is used to produce sulfuric acid. Sulfur is used to vulcanize rubber: in this case, rubber turns into rubber.
In the form of sulfur color (fine powder), sulfur is used to combat diseases of vineyards and cotton. It is used to produce gunpowder, matches, and luminous compounds. In medicine, sulfur ointments are prepared to treat skin diseases.
4.2. Hydrogen sulfide, sulfides, hydrosulfides
Hydrogen sulfide is an analogue of water. Its electronic formula
Shows that in education H-S-H bonds two p-electrons of the outer level of the sulfur atom are involved. The H 2 S molecule has an angular shape, so it is polar.
^ Being in nature
Hydrogen sulfide occurs naturally in volcanic gases and in the waters of some mineral springs, for example Pyatigorsk, Matsesta. It is formed during the decay of sulfur-containing organic substances of various animal and plant remains. This explains the characteristic bad smell Wastewater, cesspools and garbage dumps.
Receipt
1. Hydrogen sulfide can be obtained by directly combining sulfur with hydrogen when heated:
2. But it is usually 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 a Kipp apparatus.
^ Physical properties
Under normal conditions, hydrogen sulfide is a colorless gas with a strong characteristic odor. rotten eggs. Very poisonous, when inhaled it binds to hemoglobin, causing paralysis, which is often
Which leads to death. In small concentrations it is less dangerous. You need to work with him fume hoods or with hermetically sealed devices. Permissible content of H 2 S in production premises is 0.01 mg in 1 liter of air.
Hydrogen sulfide is relatively soluble in water (at 20°C, 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 hydrosulfide acid (it exhibits the properties of a weak acid).
^ Chemical properties
1, When heated strongly, hydrogen sulfide almost completely decomposes to form sulfur and hydrogen.
2. Hydrogen sulfide gas burns in air with a blue flame to form 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 of it can be explained as follows. In solution, H2S relatively easily gives up electrons to air oxygen molecules:
In this case, oxygen in the air oxidizes hydrogen sulfide to sulfur, which makes hydrogen sulfide water cloudy:
2H 2 S+O 2 =2S+2H 2 O
This also explains the fact that hydrogen sulfide does not accumulate in very large quantities in nature during the decay of organic substances - oxygen in the air oxidizes it into free sulfur.
4, Hydrogen sulfide reacts vigorously with halogen solutions, for example:
H 2 S+I 2 =2HI+S Sulfur is released and the iodine solution becomes discolored.
5. Various oxidizing agents react vigorously with hydrogen sulfide: when exposed to nitric acid free sulfur is formed.
6. Hydrogen sulfide solution has an acidic reaction due to dissociations:
H 2 SН + +HS - HS - H + +S -2
Usually the first stage predominates. It is a very weak acid: weaker than carbonic acid, which usually displaces H 2 S from sulfides.
Sulfides and hydrosulfides
Hydrogen sulfide acid, as a dibasic acid, forms two series of salts:
Medium - sulfides (Na 2 S);
Acidic - hydrosulfides (NaHS).
These salts can be obtained: - by reacting hydroxides with hydrogen sulfide: 2NaOH+H 2 S=Na 2 S+2H 2 O
Direct interaction of sulfur with metals:
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 +
Hydrosulfides are almost all highly soluble in water.
Sulfides of alkali and alkaline earth metals are also easily soluble in water and colorless.
Heavy metal sulfides 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 sulfides have an alkaline reaction when dissolved in water:
Na 2 S+ННNaHS+NaOH
All sulfides, like hydrogen sulfide itself, are energetic reducing agents:
3PbS -2 +8HN +5 O 3(diluted) =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. Based on the different solubility of sulfides and the different colors of many of them qualitative analysis cations.
^ 4.3. Sulfur(IV) oxide and sulfurous acid
Sulfur (IV) oxide, or sulfur dioxide, is under normal conditions a colorless gas with a pungent, suffocating odor. When cooled to -10°C, it liquefies into a colorless liquid.
Receipt
1. In laboratory conditions, sulfur oxide (IV) is obtained from salts of sulfurous acid by treating them with strong acids:
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 + =2SO 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 + +2SO 2- 4 =Cu 2+ + SO 2- 4 +SO 2 +2H 2 O
3. Sulfur (IV) oxide is also formed when sulfur is burned in air or oxygen:
4. Under industrial conditions, SO 2 is obtained by roasting pyrite FeS 2 or sulfur ores of non-ferrous metals (zinc blende ZnS, lead luster PbS, etc.):
4FeS 2 +11O 2 =2Fe 2 O 3 +8SO 2
Structural formula of the SO 2 molecule:
Four electrons of sulfur and four electrons from two oxygen atoms take part in the formation of bonds in a 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 (IV) oxide exhibits all the properties of acidic oxides:
Interaction with water
Interaction with alkalis,
Interaction with basic oxides.
2. Sulfur (IV) oxide 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 oxide (IV) is explained by the fact that sulfur has an oxidation state of +4 in it, and therefore it can, by donating 2 electrons, be oxidized to S +6, and by accepting 4 electrons, reduced to S°. The manifestation of these or other properties depends on the nature of the reacting component.
Sulfur oxide (IV) is highly soluble in water (40 volumes of SO 2 dissolve in 1 volume at 20°C). In this case, sulfurous acid, which exists only in an aqueous solution, is formed:
SO 2 +H 2 OH 2 SO 3
The reaction is reversible. In an aqueous solution, sulfur oxide (IV) and sulfurous acid are in chemical equilibrium, which can be shifted. When binding H 2 SO 3 (neutralization of acid
You) the reaction proceeds towards the formation of sulfurous acid; when SO 2 is removed (by blowing through a nitrogen solution or heating), the reaction proceeds towards the starting substances. A solution of sulfurous acid always contains sulfur oxide (IV), which gives it a pungent odor.
Sulfurous acid has all the properties of acids. In solution it dissociates stepwise:
H 2 SO 3 H + +HSO - 3 HSO - 3 H + +SO 2- 3
Thermally unstable, volatile. Sulfurous acid, as a dibasic acid, forms two types of salts:
Medium - sulfites (Na 2 SO 3);
Acidic - hydrosulfites (NaHSO 3).
Sulfites are formed when an acid is completely neutralized with an alkali:
H 2 SO 3 +2NaOH=Na 2 SO 3 +2H 2 O
Hydrosulfites are obtained when there is a lack of alkali:
H 2 SO 3 +NaOH=NaHSO 3 +H 2 O
Sulfurous acid and its salts have both oxidizing and reducing properties, which is determined by the nature of the reaction partner.
1. Thus, under the influence of oxygen, sulfites are oxidized to sulfates:
2Na 2 S +4 O 3 +O 0 2 =2Na 2 S +6 O -2 4
The oxidation of sulfurous acid with bromine and potassium permanganate occurs 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 energetic reducing agents, sulfites exhibit oxidizing properties:
Almost all hydrosulfites and alkali metal sulfites dissolve from sulfurous acid salts.
3. Since H 2 SO 3 is a weak acid, when acids act on sulfites and hydrosulfites, SO 2 is released. This method is usually used to obtain SO 2 in laboratory conditions:
NaHSO 3 +H 2 SO 4 =Na 2 SO 4 +SO 2 +H 2 O
4. Water-soluble sulfites are easily hydrolyzed, as a result of which the concentration of OH - ions in the solution increases:
Na 2 SO 3 +HONNaHSO 3 +NaOH
Application
Sulfur (IV) oxide and sulfurous acid decolorize many dyes, forming colorless compounds with them. The latter can decompose again when heated or exposed to light, as a result of which the color is restored. Consequently, the bleaching effect of SO 2 and H 2 SO 3 differs from the bleaching effect of chlorine. Typically, sulfur (IV) oxide is used to bleach wool, silk and straw.
Sulfur (IV) oxide kills many microorganisms. Therefore, to destroy mold fungi, they fumigate damp basements, cellars, wine barrels etc. Also used for transportation and storage of fruits and berries. Sulfur oxide IV) is used in large quantities to produce sulfuric acid.
An important application is found in a solution of calcium hydrosulfite CaHSO 3 (sulfite lye), which is used to treat wood and paper pulp.
^ 4.4. Sulfur(VI) oxide. Sulfuric acid
Sulfur oxide (VI) (see Table 20) is a colorless liquid that solidifies at a temperature of 16.8 ° C into 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 oxide (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 anhydrous acid is called oleum. Oleums can contain up to 70% SO 3 .
Receipt
1. Sulfur oxide (VI) is obtained by oxidation of sulfur dioxide with air oxygen in the presence of catalysts at a temperature of 450°C (see. Preparation of sulfuric acid):
2SO 2 +O 2 =2SO 3
2. Another way to oxidize SO 2 to SO 3 is to use nitric oxide (IV) as an oxidizing agent:
The resulting nitrogen oxide (II) when interacting with atmospheric oxygen easily and quickly turns into nitrogen oxide (IV): 2NO+O 2 = 2NO 2
Which can again be used in the oxidation of SO 2. Consequently, NO 2 acts as an oxygen carrier. This method of oxidation of SO 2 to SO 3 is called nitrous. The SO 3 molecule has the shape of a triangle, in the center of which
The sulfur atom is located:
This structure is due to the mutual repulsion of bonding electron pairs. The sulfur atom provided six outer electrons for their formation.
Chemical properties
1. SO 3 is a typical acid oxide.
2. Sulfur oxide (VI) has the properties of a strong oxidizing agent.
Application
Sulfur (VI) oxide is used to produce sulfuric acid. Highest value has a contact method of receipt
Sulfuric acid. Using this method, you can obtain 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; obtaining H 2 SO 4 .
SO 2 is obtained by roasting FeS 2 pyrite in special furnaces: 4FeS 2 +11O 2 =2Fe 2 O 3 +8SO 2
To speed up the firing, pyrite is pre-crushed, and for more complete burning of sulfur, a significant amount is introduced more air(oxygen) than required by the reaction. The gas leaving the kiln consists 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, poisons the catalyst. The gas is purified from arsenic compounds and dust by passing it through special electric filters and a washing tower; moisture is absorbed by concentrated sulfuric acid in the drying tower. Purified gas containing oxygen is heated in a heat exchanger to 450°C and enters the contact apparatus. Inside the contact apparatus there are lattice shelves filled with catalyst.
Previously, finely crushed metal platinum was used as a catalyst. Subsequently, it was replaced by vanadium compounds - vanadium (V) oxide V 2 O 5 or vanadyl sulfate VOSO 4, which are cheaper than platinum and poison more slowly.
The oxidation reaction of SO 2 to SO 3 is reversible:
2SO 2 +O 2 2SO 3
An increase in the oxygen content in the roasting gas increases the yield of sulfur oxide (VI): at a temperature of 450°C it usually reaches 95% and higher.
The resulting sulfur oxide (VI) is then fed by countercurrent into the absorption tower, where it is absorbed by concentrated sulfuric acid. As saturation occurs, anhydrous sulfuric acid is first formed, and then oleum. Subsequently, the oleum is diluted to 98% sulfuric acid and supplied to consumers.
Structural formula of sulfuric acid:
^ Physical properties
Sulfuric acid is a heavy, colorless, oily liquid that crystallizes at +10.4°C, almost twice as much ( =1.83 g/cm 3) heavier than water, odorless, non-volatile. Extremely hygroscopic. It absorbs moisture with the release of a large amount of heat, so you cannot add water to concentrated sulfuric acid - the acid will splash. For times
Add sulfuric acid to water in small portions.
Anhydrous sulfuric acid dissolves up to 70% of sulfur (VI) oxide. When heated, it splits off SO 3 until a solution with mass fraction H 2 SO 4 98.3%. Anhydrous H 2 SO 4 almost does not conduct electric current.
^ Chemical properties
1. Mixes with water in any ratio and forms hydrates of various compositions:
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 chars organic substances - sugar, paper, wood, fiber, removing water elements from them:
C 12 H 22 O 11 + H 2 SO 4 = 12 C + H 2 SO 4 11 H 2 O
The resulting carbon partially reacts with the acid:
Gas drying is based on the absorption of water by sulfuric acid.
How 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 you add H 2 SO 4 to salt solutions, then displacement of acids does not occur.
H 2 SO 4 is a 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.
Dilute sulfuric acid is characterized by all the properties of non-oxidizing acids. Namely: it interacts with metals that are in the electrochemical series of metal voltages up to hydrogen:
Interaction with metals occurs due to the reduction of hydrogen ions.
6. Concentrated sulfuric acid is a vigorous oxidizing agent. When heated, it oxidizes most metals, including those in the electrochemical voltage series after hydrogen. It 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 strong metals such as aluminum, iron, and chromium. This is explained by the passivation of metals. This feature is widely used when transporting it in iron containers.
However, when heated:
Thus, concentrated sulfuric acid interacts with metals due to the reduction of acid-forming atoms.
The 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 an essential product of the basic chemical industry involved in the production of non-
Organic acids, alkalis, salts, mineral fertilizers and chlorine.
In terms of variety of applications, sulfuric acid ranks first among acids. Largest quantity it is used to produce phosphorus and nitrogen fertilizers. Being non-volatile, sulfuric acid is used to produce other acids - hydrochloric, hydrofluoric, phosphoric and acetic.
A lot of it is used to purify petroleum products - gasoline, kerosene, lubricating oils - from harmful impurities. In mechanical engineering, sulfuric acid is used to clean the metal surface from oxides before coating (nickel plating, chrome plating, etc.). Sulfuric acid is used in the production of explosives, artificial fibers, dyes, plastics and many others. It is used to fill batteries.
Salts of sulfuric acid are important.
^ Sodium sulfate Na 2 SO 4 crystallizes from aqueous solutions in the form of Na 2 SO 4 10H 2 O hydrate, which is called Glauber's salt. Used in medicine as a laxative. Anhydrous sodium sulfate is used in the production of soda and glass.
^ Ammonium sulfate(NH 4) 2 SO 4 - nitrogen fertilizer.
Potassium sulfate K 2 SO 4 - potassium fertilizer.
Calcium sulfate CaSO 4 occurs in nature in the form of the gypsum mineral CaSO 4 2H 2 O. When heated to 150°C, it loses part of the water and turns into a hydrate of the composition 2CaSO 4 H 2 O, called burnt gypsum, or alabaster. Alabaster, when mixed with water into a dough-like mass, after some time hardens again, turning into CaSO 4 2H 2 O. Gypsum is widely used in construction (plaster).
^ Magnesium sulfate MgSO 4 is contained in sea water, causing its bitter taste. Crystal hydrate, called bitter salt, is used as a laxative.
Vitriol- technical name for crystalline hydrates of metal sulfates Fe, Cu, Zn, Ni, Co (dehydrated salts are not vitriol). Copper sulfate CuSO 4 5H 2 O - toxic substance of blue color. Its diluted solution is sprayed on plants and the seeds are treated before sowing. inkstone FeSO 4 7H 2 O is a light green substance. Used to control plant pests, prepare inks, mineral paints, etc. Zinc sulfate ZnSO 4 7H 2 O is used in the production of mineral paints, in calico printing, and medicine.
^ 4.5. Esters of sulfuric acid. Sodium thiosulfate
Esters of sulfuric acid include dialkyl sulfates (RO 2)SO 2. These are high-boiling liquids; lower ones are soluble in water; in the presence of alkalis they form alcohol and sulfuric acid salts. Lower dialkyl sulfates are alkylating agents.
Diethyl sulfate(C 2 H 5) 2 SO 4. Melting point -26°C, boiling point 210°C, soluble in alcohols, insoluble in water. Obtained by reacting sulfuric acid with ethanol. It is an ethylation agent in organic synthesis. Penetrates through the skin.
Dimethyl sulfate(CH 3) 2 SO 4. Melting point -26.8°C, boiling point 188.5°C. Soluble in alcohols, poorly soluble in water. Reacts with ammonia in the absence of a solvent (explosively); sulfonates some aromatic compounds, such as phenol esters. It is obtained by reacting 60% oleum with methanol at 150°C. It is a methylating agent in organic synthesis. Carcinogen, affects eyes, skin, respiratory organs.
^ Sodium thiosulfate Na2S2O3
A salt of thiosulfuric acid in which the two sulfur atoms have different oxidation states: +6 and -2. Crystalline substance, highly soluble in water. It is produced in the form of crystalline hydrate Na 2 S 2 O 3 5H 2 O, commonly called hyposulfite. It is obtained by reacting 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 to absorb chlorine (in the first gas masks) was based on this reaction.
The oxidation of sodium thiosulfate by weak oxidizing agents occurs somewhat differently. In this case, salts of tetrathionic acid are formed, 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, during the purification of industrial gases from sulfur. Used to remove traces of chlorine after bleaching fabrics, to extract silver from ores; It is a fixative in photography, a reagent in iodometry, an antidote for poisoning with arsenic and mercury compounds, and an anti-inflammatory agent.
Sulfur dioxide is a colorless gas with a pungent odor. The molecule has an angular shape.
- Melting point - -75.46 °C,
- Boiling point - -10.6 °C,
- Gas density - 2.92655 g/l.
Easily liquefies into a colorless, highly mobile liquid at a temperature of 25 ° C and a pressure of about 0.5 MPa.
For liquid form the density is 1.4619 g/cm 3 (at - 10 °C).
Solid sulfur dioxide - colorless crystals, orthorhombic system.
Sulfur dioxide dissociates noticeably only around 2800 °C.
Dissociation of liquid sulfur dioxide proceeds according to the following scheme:
2SO 2 ↔ SO 2+ + SO 3 2-
Three-dimensional model of a molecule
The solubility of sulfur dioxide in water depends on temperature:
- at 0 °C 22.8 g of sulfur dioxide dissolves in 100 g of water,
- at 20 °C - 11.5 g,
- at 90 °C - 2.1 g.
An aqueous solution of sulfur dioxide is sulfurous acid H 2 SO 3.
Sulfur dioxide is soluble in ethanol, H 2 SO 4, oleum, CH 3 COOH. Liquid sulfur dioxide is mixed in any ratio with SO 3. CHCl 3, CS 2, diethyl ether.
Liquid sulfur dioxide dissolves chlorides. Metal iodides and thiocyanates do not dissolve.
Salts dissolved in liquid sulfur dioxide dissociate.
Sulfur dioxide is capable of being reduced to sulfur and oxidized to hexavalent sulfur compounds.
Sulfur dioxide is toxic. At a concentration of 0.03-0.05 mg/l, it irritates the mucous membranes, respiratory organs, and eyes.
The main industrial method for producing sulfur dioxide is from sulfur pyrite FeS 2 by burning it and further processing with weak cold H 2 SO 4.
In addition, sulfur dioxide can be produced by burning sulfur, and also as a by-product of roasting copper and zinc sulfide ores.
Sulfide sulfur is available to plants only after converting to the sulfate form. Most of the sulfur is present in the soil as organic compounds, not absorbed by plants. Only after the mineralization of organic substances and the transition of sulfur to the sulfate form does organic sulfur become available to plants.
Chemical industry does not produce fertilizers with basic active substance sulfur dioxide. However, it is found as an impurity in many fertilizers. These include phosphogypsum, simple superphosphate, ammonium sulfate, potassium sulfate, potassium magnesia, gypsum, oil shale ash, manure, peat and many others.
Absorption of sulfur dioxide by plants
Sulfur enters plants through the roots in the form SO 4 2- and leaves in the form of sulfur dioxide. At the same time, the absorption of sulfur from the atmosphere provides up to 80% of the plants’ needs for this element. In this regard, near industrial centers, where the atmosphere is rich in sulfur dioxide, plants are well supplied with sulfur. In remote areas, the amount of sulfur dioxide in precipitation and the atmosphere is greatly reduced and the nutrition of plants with sulfur depends on its presence in the soil.