Sulfur oxide 4 plus chlorine. Sulfur(IV) oxide and sulfurous acid
Hydrogen sulfide – H2S
Sulfur compounds -2, +4, +6. Qualitative reactions to sulfides, sulfites, sulfates.
Receipt upon interaction:
1. hydrogen with sulfur at t – 300 0
2. when acting on sulfides of mineral acids:
Na 2 S+2HCl =2 NaCl+H 2 S
Physical properties:
a colorless gas with the smell of rotten eggs, poisonous, heavier than air, and dissolving in water to form weak hydrogen sulfide acid.
Chemical properties
Acid-base properties
1. A solution of hydrogen sulfide in water - hydrosulfide acid - is a weak dibasic acid, therefore it dissociates stepwise:
H 2 S ↔ HS - + H +
HS - ↔ H - + S 2-
2. Hydrogen sulfide acid has the general properties of acids, reacts with metals, basic oxides, bases, salts:
H 2 S + Ca = CaS + H 2
H 2 S + CaO = CaS + H 2 O
H 2 S + 2NaOH = Na 2 S + 2H 2 O
H 2 S + CuSO 4 = CuS↓ + H 2 SO 4
All acid salts - hydrosulfides - are highly soluble in water. Normal salts - sulfides - dissolve in water in different ways: sulfides of alkali and alkaline earth metals are highly soluble, sulfides of other metals are insoluble in water, and sulfides of copper, lead, mercury and some other heavy metals are not soluble even in acids (except nitric acid)
CuS+4HNO 3 =Cu(NO 3) 2 +3S+2NO+2H 2 O
Soluble sulfides undergo hydrolysis - at the anion.
Na 2 S ↔ 2Na + + S 2-
S 2- +HOH ↔HS - +OH -
Na 2 S + H 2 O ↔ NaHS + NaOH
A qualitative reaction to hydrosulfide acid and its soluble salts (i.e., to the sulfide ion S 2-) is their interaction with soluble lead salts, which results in the formation of a black PbS precipitate
Na 2 S + Pb(NO 3) 2 = 2NaNO 3 + PbS↓
Pb 2+ + S 2- = PbS↓
Shows only restorative properties, because the sulfur atom has the lowest oxidation state -2
1. with oxygen
a) with a disadvantage
2H 2 S -2 +O 2 0 = S 0 +2H 2 O -2
b) with excess oxygen
2H 2 S+3O 2 =2SO 2 +2H 2 O
2. with halogens (bromine water discoloration)
H 2 S -2 +Br 2 =S 0 +2HBr -1
3. with conc. HNO3
H 2 S+2HNO 3 (k) = S+2NO 2 +2H 2 O
b) with strong oxidizing agents (KMnO 4, K 2 CrO 4 in an acidic environment)
2KMnO 4 +3H 2 SO 4 +5H 2 S = 5S+2MnSO 4 +K 2 SO 4 +8H 2 O
c) hydrosulfide acid is oxidized not only by strong oxidizing agents, but also by weaker ones, for example, iron (III) salts, sulfurous acid, etc.
2FeCl 3 + H 2 S = 2FeCl 2 + S + 2HCl
H 2 SO 3 + 2H 2 S = 3S + 3H 2 O
Receipt
1. combustion of sulfur in oxygen.
2. combustion of hydrogen sulfide in excess O 2
2H 2 S+3O 2 = 2SO 2 +2H 2 O
3. sulfide oxidation
2CuS+3O2 = 2SO2 +2CuO
4. interaction of sulfites with acids
Na 2 SO 3 +H 2 SO 4 =Na 2 SO 4 +SO 2 +H 2 O
5. interaction of metals in the activity series after (H 2) with conc. H2SO4
Cu+2H 2 SO 4 = CuSO 4 + SO 2 +2H 2 O
Physical properties
Gas, colorless, with a suffocating odor of burnt sulfur, poisonous, more than 2 times heavier than air, highly soluble in water (at room temperature, about 40 volumes of gas dissolve in one volume).
Chemical properties:
Acid-base properties
SO 2 is a typical acidic oxide.
1.with alkalis, forming two types of salts: sulfites and hydrosulfites
2KOH+SO2 = K2SO3 +H2O
KOH+SO 2 = KHSO 3 +H 2 O
2.with basic oxides
K 2 O+SO 2 = K 2 SO 3
3. weak sulfurous acid is formed with water
H 2 O+SO 2 = H 2 SO 3
Sulfurous acid exists only in solution and is a weak acid.
has all the general properties of acids.
4. qualitative reaction to sulfite - ion - SO 3 2 - action of mineral acids
Na 2 SO 3 +2HCl= 2Na 2 Cl+SO 2 +H 2 O smell of burnt sulfur
Redox properties
In ORR it can be both an oxidizing agent and a reducing agent, because the sulfur atom in SO 2 has an intermediate oxidation state of +4.
As an oxidizing agent:
SO 2 + 2H 2 S = 3S + 2H 2 S
As a reducing agent:
2SO 2 +O 2 = 2SO 3
Cl 2 +SO 2 +2H 2 O = H 2 SO 4 +2HCl
2KMnO 4 +5SO 2 +2H 2 O = K 2 SO 4 +2H 2 SO 4 +2MnSO 4
Sulfur oxide (VI) SO 3 (sulfuric anhydride)
Receipt:
Oxidation of sulfur dioxide
2SO 2 + O 2 = 2SO 3 ( t 0 , kat)
Physical properties
A colorless liquid, at temperatures below 17 0 C it turns into a white crystalline mass. Thermally unstable compound, completely decomposes at 700 0 C. It is highly soluble in water and anhydrous sulfuric acid and reacts with it to form oleum
SO 3 + H 2 SO 4 = H 2 S 2 O 7
Chemical properties
Acid-base properties
Typical acid oxide.
1.with alkalis, forming two types of salts: sulfates and hydrosulfates
2KOH+SO 3 = K 2 SO 4 +H 2 O
KOH+SO 3 = KHSO 4 +H 2 O
2.with basic oxides
CaO+SO 2 = CaSO 4
3. with water
H 2 O + SO 3 = H 2 SO 4
Redox properties
Sulfur oxide (VI) is a strong oxidizing agent, usually reduced to SO 2
3SO 3 + H 2 S = 4SO 2 + H 2 O
Sulfuric acid H 2 SO 4
Preparation of sulfuric acid
In industry, acid is produced by contact method:
1. pyrite firing
4FeS 2 +11O 2 = 2Fe 2 O 3 + 8SO 2
2. oxidation of SO 2 to SO 3
2SO 2 + O 2 = 2SO 3 ( t 0 , kat)
3. dissolution of SO 3 in sulfuric acid
n SO 3 + H 2 SO 4 = H 2 SO 4 ∙ n SO 3 (oleum)
H2SO4∙ n SO 3 + H 2 O = H 2 SO 4
Physical properties
H 2 SO 4 is a heavy oily liquid, odorless and colorless, hygroscopic. Mixes with water in any ratio; when concentrated sulfuric acid is dissolved in water, a large amount of heat is released, so it must be carefully poured into water, and not vice versa (first water, then acid, otherwise big trouble will happen)
A solution of sulfuric acid in water with a H 2 SO 4 content of less than 70% is usually called dilute sulfuric acid, more than 70% - concentrated.
Chemical properties
Acid-base
Dilute sulfuric acid exhibits all the characteristic properties of strong acids. Dissociates in aqueous solution:
H 2 SO 4 ↔ 2H + + SO 4 2-
1. with basic oxides
MgO + H 2 SO 4 = MgSO 4 + H 2 O
2. with grounds
2NaOH +H 2 SO 4 = Na 2 SO 4 + 2H 2 O
3. with salts
BaCl 2 + H 2 SO 4 = BaSO 4 ↓ + 2HCl
Ba 2+ + SO 4 2- = BaSO 4 ↓ (white precipitate)
Qualitative reaction to sulfate ion SO 4 2-
Due to its higher boiling point, compared to other acids, sulfuric acid, when heated, displaces them from salts:
NaCl + H 2 SO 4 = HCl + NaHSO 4
Redox properties
In dilute H 2 SO 4 the oxidizing agents are H + ions, and in concentrated H 2 SO 4 the oxidizing agents are SO 4 2 sulfate ions.
Metals in the activity series up to hydrogen dissolve in dilute sulfuric acid, sulfates are formed and hydrogen is released
Zn + H 2 SO 4 = ZnSO 4 + H 2
Concentrated sulfuric acid is a vigorous oxidizing agent, especially when heated. It oxidizes many metals, non-metals, inorganic and organic substances.
H 2 SO 4 (k) oxidizing agent S +6
With more active metals, sulfuric acid, depending on the concentration, can be reduced to a variety of products
Zn + 2H 2 SO 4 = ZnSO 4 + SO 2 + 2H 2 O
3Zn + 4H 2 SO 4 = 3ZnSO 4 + S + 4H 2 O
4Zn + 5H 2 SO 4 = 4ZnSO 4 + H 2 S + 4H 2 O
Concentrated sulfuric acid oxidizes some non-metals (sulfur, carbon, phosphorus, etc.), reducing to sulfur oxide (IV)
S + 2H 2 SO 4 = 3SO 2 + 2H 2 O
C + 2H 2 SO 4 = 2SO 2 + CO 2 + 2H 2 O
Interaction with some complex substances
H 2 SO 4 + 8HI = 4I 2 + H 2 S + 4 H 2 O
H 2 SO 4 + 2HBr = Br 2 + SO 2 + 2H 2 O
Sulfuric acid salts
2 types of salts: sulfates and hydrosulfates
Salts of sulfuric acid have all the general properties of salts. Their relationship to heat is special. Sulfates of active metals (Na, K, Ba) do not decompose even when heated above 1000 0 C, salts of less active metals (Al, Fe, Cu) decompose even with slight heating
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 compound in the 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 the chemical formula SO 2. 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 latter characteristic is 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 indicating poisoning are coughing, breathing problems, usually in the form of a runny nose, hoarseness, 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/m 3 . 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.
Physical qualitative 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, and 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 types of SO 3 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 the environment during the use of 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 valency. May be highly toxic and cause serious health problems.
Sulfur dioxide has a molecular structure similar to ozone. The sulfur atom at the center of the molecule is bonded to two oxygen atoms. This gaseous product of sulfur oxidation is colorless, emits a pungent odor, and easily condenses into a clear liquid when conditions change. The substance is highly soluble in water and has antiseptic properties. SO 2 is produced in large quantities in the chemical industry, namely in the sulfuric acid production cycle. The gas is widely used for processing agricultural and food products, bleaching fabrics in the textile industry.
Systematic and trivial names of substances
It is necessary to understand the variety of terms related to the same compound. The official name of the compound, the chemical composition of which is reflected by the formula SO 2, is sulfur dioxide. IUPAC recommends using this term and its English equivalent - Sulfur dioxide. Textbooks for schools and universities often mention another name - sulfur (IV) oxide. The Roman numeral in parentheses indicates the valency of the S atom. Oxygen in this oxide is divalent, and the oxidation number of sulfur is +4. In the technical literature, outdated terms such as sulfur dioxide, sulfuric acid anhydride (a product of its dehydration) are used.
Composition and features of the molecular structure of SO 2
The SO 2 molecule is formed by one sulfur atom and two oxygen atoms. There is an angle of 120° between covalent bonds. In the sulfur atom, sp2 hybridization occurs—the clouds of one s and two p electrons are aligned in shape and energy. They are the ones who participate in the formation of a covalent bond between sulfur and oxygen. In the O–S pair, the distance between the atoms is 0.143 nm. Oxygen is a more electronegative element than sulfur, which means that the bonding pairs of electrons shift from the center to the outer corners. The entire molecule is also polarized, the negative pole is the O atoms, the positive pole is the S atom.
Some physical parameters of sulfur dioxide
Quadrivalent sulfur oxide, under normal environmental conditions, retains a gaseous state of aggregation. The formula of sulfur dioxide allows you to determine its relative molecular and molar mass: Mr(SO 2) = 64.066, M = 64.066 g/mol (can be rounded to 64 g/mol). This gas is almost 2.3 times heavier than air (M(air) = 29 g/mol). Dioxide has a sharp, specific smell of burning sulfur, which is difficult to confuse with any other. It is unpleasant, irritates the mucous membranes of the eyes, and causes a cough. But sulfur (IV) oxide is not as poisonous as hydrogen sulfide.
Under pressure at room temperature, sulfur dioxide gas liquefies. At low temperatures, the substance is in a solid state and melts at -72...-75.5 °C. With a further increase in temperature, liquid appears, and at -10.1 °C gas is formed again. SO 2 molecules are thermally stable; decomposition into atomic sulfur and molecular oxygen occurs at very high temperatures (about 2800 ºC).
Solubility and interaction with water
Sulfur dioxide, when dissolved in water, partially reacts with it to form a very weak sulfurous acid. At the moment of receipt, it immediately decomposes into anhydride and water: SO 2 + H 2 O ↔ H 2 SO 3. In fact, it is not sulfurous acid that is present in the solution, but hydrated SO 2 molecules. Dioxide gas reacts better with cool water, and its solubility decreases with increasing temperature. Under normal conditions, up to 40 volumes of gas can dissolve in 1 volume of water.
Sulfur dioxide in nature
Significant volumes of sulfur dioxide are released with volcanic gases and lava during eruptions. Many types of anthropogenic activities also lead to increased concentrations of SO 2 in the atmosphere.
Sulfur dioxide is released into the air by metallurgical plants, where waste gases are not captured during ore roasting. Many types of fossil fuels contain sulfur; as a result, significant volumes of sulfur dioxide are released into the atmospheric air when burning coal, oil, gas, and fuel obtained from them. Sulfur dioxide becomes toxic to humans at concentrations in the air above 0.03%. A person begins to experience shortness of breath, and symptoms resembling bronchitis and pneumonia may occur. Very high concentrations of sulfur dioxide in the atmosphere can lead to severe poisoning or death.
Sulfur dioxide - production in the laboratory and in industry
Laboratory methods:
- When sulfur is burned in a flask with oxygen or air, dioxide is obtained according to the formula: S + O 2 = SO 2.
- You can act on the salts of sulfurous acid with stronger inorganic acids, it is better to take hydrochloric acid, but you can use diluted sulfuric acid:
- Na 2 SO 3 + 2HCl = 2NaCl + H 2 SO 3;
- Na 2 SO 3 + H 2 SO 4 (diluted) = Na 2 SO 4 + H 2 SO 3;
- H 2 SO 3 = H 2 O + SO 2.
3. When copper reacts with concentrated sulfuric acid, it is not hydrogen that is released, but sulfur dioxide:
2H 2 SO 4 (conc.) + Cu = CuSO 4 + 2H 2 O + SO 2.
Modern methods of industrial production of sulfur dioxide:
- Oxidation of natural sulfur when it is burned in special furnaces: S + O 2 = SO 2.
- Firing iron pyrite (pyrite).
Basic chemical properties of sulfur dioxide
Sulfur dioxide is a chemically active compound. In redox processes, this substance often acts as a reducing agent. For example, when molecular bromine reacts with sulfur dioxide, the reaction products are sulfuric acid and hydrogen bromide. The oxidizing properties of SO 2 appear if this gas is passed through hydrogen sulfide water. As a result, sulfur is released, self-oxidation-self-reduction occurs: SO 2 + 2H 2 S = 3S + 2H 2 O.
Sulfur dioxide exhibits acidic properties. It corresponds to one of the weakest and most unstable acids - sulfurous. This compound does not exist in its pure form; the acidic properties of a sulfur dioxide solution can be detected using indicators (litmus turns pink). Sulfurous acid produces medium salts - sulfites and acidic salts - hydrosulfites. Among them there are stable compounds.
The process of oxidation of sulfur in dioxide to the hexavalent state in sulfuric acid anhydride is catalytic. The resulting substance dissolves vigorously in water and reacts with H 2 O molecules. The reaction is exothermic, and sulfuric acid is formed, or rather, its hydrated form.
Practical uses of sulfur dioxide
The main method of industrial production of sulfuric acid, which requires elemental dioxide, has four stages:
- Obtaining sulfur dioxide by burning sulfur in special furnaces.
- Purification of the resulting sulfur dioxide from all kinds of impurities.
- Further oxidation to hexavalent sulfur in the presence of a catalyst.
- Absorption of sulfur trioxide by water.
Previously, almost all of the sulfur dioxide needed to produce sulfuric acid on an industrial scale was obtained by roasting pyrite as a by-product of steelmaking. New types of processing of metallurgical raw materials use less ore combustion. Therefore, natural sulfur has become the main starting material for sulfuric acid production in recent years. Significant global reserves of this raw material and its availability make it possible to organize large-scale processing.
Sulfur dioxide is widely used not only in the chemical industry, but also in other sectors of the economy. Textile mills use this substance and the products of its chemical reaction to bleach silk and wool fabrics. This is a type of chlorine-free bleaching that does not destroy the fibers.
Sulfur dioxide has excellent disinfectant properties, which is used in the fight against fungi and bacteria. Sulfur dioxide is used to fumigate agricultural storage facilities, wine barrels and cellars. SO 2 is used in the food industry as a preservative and antibacterial substance. They add it to syrups and soak fresh fruits in it. Sulfitization
Sugar beet juice decolorizes and disinfects raw materials. Canned vegetable purees and juices also contain sulfur dioxide as an antioxidant and preservative.
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 resonance 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 on 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 the technological conditions for fermentation and storage of wine.
These conditions, with the correct use of 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 released by metallurgical plants (source: coking coal), as well as a number of chemical industries (for example, sulfuric acid production). 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 the bleaching of 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, carbon dioxide is formed and released in the air of working rooms in a number of industries: sulfuric acid, cellulose, during roasting of ores containing sulfur metals, in pickling rooms at metal factories, in the production of glass, ultramarine, etc., very often sulfur is contained in the air of boiler rooms and ash rooms , where it is formed by burning 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
Sulfur is widespread in the earth's crust and ranks sixteenth among other elements. It is found both in a free state and in a bound form. Non-metallic properties are characteristic of this chemical element. Its Latin name is "Sulfur", denoted by the symbol S. The element is part of various ion compounds containing oxygen and/or hydrogen, forms many substances belonging to the classes of acids, salts and several oxides, each of which can be called sulfur oxide with the addition symbols indicating valence. The oxidation states that it exhibits in various compounds are +6, +4, +2, 0, −1, −2. Sulfur oxides with varying degrees of oxidation are known. The most common are sulfur dioxide and trioxide. Less known are sulfur monoxide, as well as higher (except SO3) and lower oxides of this element.
Sulfur monoxide
An inorganic compound called sulfur oxide II, SO, is a colorless gas in appearance. Upon contact with water, it does not dissolve, but reacts with it. This is a very rare compound that is found only in a rarefied gas environment. The SO molecule is thermodynamically unstable and initially turns into S2O2 (called disulfur gas or sulfur peroxide). Due to the rare occurrence of sulfur monoxide in our atmosphere and the low stability of the molecule, it is difficult to fully determine the dangers of this substance. But in condensed or more concentrated form, the oxide turns into peroxide, which is relatively toxic and caustic. This compound is also highly flammable (reminiscent of methane in this property); when burned, it produces sulfur dioxide, a poisonous gas. Sulfur oxide 2 was discovered near Io (one of the atmospheres of Venus and the interstellar medium. On Io it is believed to be produced by volcanic and photochemical processes. The main photochemical reactions are as follows: O + S2 → S + SO and SO2 → SO + O.
Sulphur dioxide
Sulfur oxide IV, or sulfur dioxide (SO2), is a colorless gas with a suffocating, pungent odor. At a temperature of minus 10 C it turns into a liquid state, and at a temperature of minus 73 C it solidifies. At 20C, about 40 volumes of SO2 dissolve in 1 liter of water.
This sulfur oxide, dissolving in water, forms sulfurous acid, since it is its anhydride: SO2 + H2O ↔ H2SO3.
It interacts with bases and 2NaOH + SO2 → Na2SO3 + H2O and SO2 + CaO → CaSO3.
Sulfur dioxide is characterized by the properties of both an oxidizing agent and a reducing agent. It is oxidized by atmospheric oxygen to sulfuric anhydride in the presence of a catalyst: SO2 + O2 → 2SO3. With strong reducing agents such as hydrogen sulfide, it plays the role of an oxidizing agent: H2S + SO2 → S + H2O.
Sulfur dioxide is used in industry mainly to produce sulfuric acid. Sulfur dioxide is produced by burning sulfur or iron pyrites: 11O2 + 4FeS2 → 2Fe2O3 + 8SO2.
Sulfuric anhydride
Sulfur oxide VI, or sulfur trioxide (SO3) is an intermediate product and has no independent significance. In appearance it is a colorless liquid. It boils at a temperature of 45 C, and below 17 C it turns into a white crystalline mass. This sulfur (with the oxidation state of the sulfur atom + 6) is extremely hygroscopic. With water it forms sulfuric acid: SO3 + H2O ↔ H2SO4. When dissolved in water, it releases a large amount of heat and, if a large amount of oxide is added not gradually, but immediately, an explosion can occur. Sulfur trioxide dissolves well in concentrated sulfuric acid to form oleum. The SO3 content in oleum reaches 60%. This sulfur compound has all the properties
Higher and lower sulfur oxides
Sulfurs are a group of chemical compounds with the formula SO3 + x, where x can be 0 or 1. The monomeric oxide SO4 contains a peroxo group (O-O) and is characterized, like the oxide SO3, by the oxidation state of sulfur +6. This sulfur oxide can be produced at low temperatures (below 78 K) from the reaction of SO3 and or photolysis of SO3 mixed with ozone.
Lower sulfur oxides are a group of chemical compounds that include:
- SO (sulfur oxide and its dimer S2O2);
- sulfur monoxides SnO (are cyclic compounds consisting of rings formed by sulfur atoms, and n can be from 5 to 10);
- S7O2;
- polymer sulfur oxides.
Interest in lower sulfur oxides has increased. This is due to the need to study their content in terrestrial and extraterrestrial atmospheres.
