In redox processes, sulfur dioxide can be both an oxidizing agent and a reducing agent because the atom in this compound has an intermediate oxidation state of +4.

How SO 2 reacts with stronger reducing agents, such as:

SO 2 + 2H 2 S = 3S↓ + 2H 2 O

How does the reducing agent SO 2 react with stronger oxidizing agents, for example with in the presence of a catalyst, with, etc.:

2SO2 + O2 = 2SO3

SO 2 + Cl 2 + 2H 2 O = H 2 SO 3 + 2HCl

Receipt

1) Sulfur dioxide is formed when sulfur burns:

2) In industry it is obtained by roasting pyrite:

3) In the laboratory, sulfur dioxide can be obtained:

Cu + 2H 2 SO 4 = CuSO 4 + SO 2 + 2H 2 O

Application

Sulfur dioxide finds wide application V textile industry for bleaching various products. In addition, it is used in agriculture for the destruction of harmful microorganisms in greenhouses and cellars. IN large quantities SO 2 is used to produce sulfuric acid.

Sulfur oxide (VI) – SO 3 (sulfuric anhydride)

Sulfuric anhydride SO 3 is a colorless liquid, which at temperatures below 17 o C turns into a white crystalline mass. Absorbs moisture very well (hygroscopic).

Chemical properties

Acid-base properties

How a typical acid oxide, sulfuric anhydride, reacts:

SO 3 + CaO = CaSO 4

c) with water:

SO 3 + H 2 O = H 2 SO 4

A special property of SO 3 is its ability to dissolve well in sulfuric acid. A solution of SO 3 in sulfuric acid is called oleum.

Formation of oleum: H 2 SO 4 + n SO 3 = H 2 SO 4 ∙ n SO 3

Redox properties

Sulfur oxide (VI) is characterized by strong oxidizing properties (usually reduced to SO 2):

3SO 3 + H 2 S = 4SO 2 + H 2 O

Receipt and use

Sulfuric anhydride is formed by the oxidation of sulfur dioxide:

2SO2 + O2 = 2SO3

IN pure form sulfuric anhydride practical significance does not have. It is obtained as an intermediate product in the production of sulfuric acid.

H2SO4

Mention of sulfuric acid was first found among Arab and European alchemists. It was obtained by calcining iron sulfate (FeSO 4 ∙ 7H 2 O) in air: 2FeSO 4 = Fe 2 O 3 + SO 3 + SO 2 or a mixture with: 6KNO 3 + 5S = 3K 2 SO 4 + 2SO 3 + 3N 2, and the released sulfuric anhydride vapors condensed. Absorbing moisture, they turned into oleum. Depending on the method of preparation, H 2 SO 4 was called oil of vitriol or sulfur oil. In 1595, the alchemist Andreas Liebavius ​​established the identity of both substances.

For a long time, oil of vitriol was not widely used. Interest in it increased greatly after in the 18th century. The process of obtaining indigo carmine, a stable blue dye, from indigo was discovered. The first factory for the production of sulfuric acid was founded near London in 1736. The process was carried out in lead chambers, at the bottom of which water was poured. A molten mixture of saltpeter and sulfur was burned in the upper part of the chamber, then air was introduced into it. The procedure was repeated until an acid of the required concentration was formed at the bottom of the container.

In the 19th century the method was improved: instead of saltpeter they began to use nitric acid(it gives when decomposed in the chamber). To return nitrous gases to the system, special towers were constructed, which gave the name to the whole process - the tower process. Factories operating using the tower method still exist today.

Sulfuric acid– it is a heavy oily liquid, colorless and odorless, hygroscopic; dissolves well in water. When concentrated sulfuric acid is dissolved in water, a large amount of heat is released, so it must be carefully poured into the water (and not vice versa!) and the solution must be mixed.

A solution of sulfuric acid in water with a H 2 SO 4 content of less than 70% is usually called dilute sulfuric acid, and a solution of more than 70% is concentrated sulfuric acid.

Chemical properties

Acid-base properties

Dilute sulfuric acid reveals everything characteristic properties strong acids. She reacts:

H 2 SO 4 + NaOH = Na 2 SO 4 + 2H 2 O

H 2 SO 4 + BaCl 2 = BaSO 4 ↓ + 2HCl

The process of interaction of Ba 2+ ions with SO 4 2+ sulfate ions leads to the formation of a white insoluble precipitate BaSO 4 . This qualitative reaction to sulfate ion.

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. SO 4 2+ ions are stronger oxidizing agents than H + ions (see diagram).

IN dilute sulfuric acid metals that are in the electrochemical voltage series are dissolved to hydrogen. In this case, metal sulfates are formed and the following is released:

Zn + H 2 SO 4 = ZnSO 4 + H 2

Metals that are located after hydrogen in the electrochemical voltage series do not react with dilute sulfuric acid:

Cu + H 2 SO 4 ≠

Concentrated sulfuric acid is a strong oxidizing agent, especially when heated. It oxidizes many and some organic substances.

When concentrated sulfuric acid interacts with metals that are located after hydrogen in the electrochemical voltage series (Cu, Ag, Hg), metal sulfates are formed, as well as the reduction product of sulfuric acid - SO 2.

Reaction of sulfuric acid with zinc

With more active metals (Zn, Al, Mg), concentrated sulfuric acid can be reduced to free sulfuric acid. For example, when sulfuric acid reacts with, depending on the concentration of the acid, various reduction products of sulfuric acid - SO 2, S, H 2 S - can be formed simultaneously:

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

In the cold, concentrated sulfuric acid passivates some metals, for example and, so it is transported in iron tanks:

Fe + H 2 SO 4 ≠

Concentrated sulfuric acid oxidizes some non-metals (, etc.), reducing to sulfur oxide (IV) SO 2:

S + 2H 2 SO 4 = 3SO 2 + 2H 2 O

C + 2H 2 SO 4 = 2SO 2 + CO 2 + 2H 2 O

Receipt and use

In industry, sulfuric acid is produced by contact method. The obtaining process occurs in three stages:

1. Obtaining SO 2 by roasting pyrite:

4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2

1. Oxidation of SO 2 to SO 3 in the presence of a catalyst – vanadium (V) oxide:

2SO2 + O2 = 2SO3

1. Dissolution of SO 3 in sulfuric acid:

H2SO4+ n SO 3 = H 2 SO 4 ∙ n SO 3

The resulting oleum is transported in iron tanks. Sulfuric acid of the required concentration is obtained from oleum by adding it to water. This can be expressed by the diagram:

H2SO4∙ n SO 3 + H 2 O = H 2 SO 4

Sulfuric acid has a variety of uses in a wide range of applications national economy. It is used for drying gases, in the production of other acids, for the production of fertilizers, various dyes and medicines.

Sulfuric acid salts

Most sulfates are highly soluble in water (CaSO 4 is slightly soluble, PbSO 4 is even less soluble and BaSO 4 is practically insoluble). Some sulfates containing water of crystallization are called vitriols:

CuSO 4 ∙ 5H 2 O copper sulfate

FeSO 4 ∙ 7H 2 O iron sulfate

Everyone has salts of sulfuric acid. Their relationship to heat is special.

Sulfates of active metals (,) do not decompose even at 1000 o C, while others (Cu, Al, Fe) decompose with slight heating into metal oxide and SO 3:

CuSO 4 = CuO + SO 3

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*in the recording image is a photograph of copper sulfate

DEFINITION

Anhydrous sulfuric acid is a heavy, viscous liquid that is easily miscible with water in any proportion: the interaction is characterized by an extremely large exothermic effect (~880 kJ/mol at infinite dilution) and can lead to explosive boiling and splashing of the mixture if water is added to the acid; This is why it is so important to always reverse the order in preparing solutions and add the acid to the water, slowly and with stirring.

Some physical properties of sulfuric acid are given in the table.

Anhydrous H 2 SO 4 is a remarkable compound with unusually high dielectric constant and very high electrical conductivity, which is due to ionic autodissociation (autoprotolysis) of the compound, as well as the relay mechanism of conductivity with proton transfer, which ensures the occurrence of electric current through a viscous liquid with a large number hydrogen bonds.

Table 1. Physical properties sulfuric acid.

Preparation of sulfuric acid

Sulfuric acid is the most important industrial chemical and the cheapest acid produced in large volume anywhere in the world.

Concentrated sulfuric acid (“oil of vitriol”) was first obtained by heating “green vitriol” FeSO 4 ×nH 2 O and consumed in large quantities to obtain Na 2 SO 4 and NaCl.

IN modern process To produce sulfuric acid, a catalyst consisting of vanadium(V) oxide with the addition of potassium sulfate on a carrier of silicon dioxide or kieselguhr is used. Sulfur dioxide SO2 is produced by burning pure sulfur or by roasting sulfide ore (primarily pyrite or ores of Cu, Ni and Zn) in the process of extracting these metals. SO2 is then oxidized to trioxide, and then sulfuric acid is obtained by dissolving in water:

S + O 2 → SO 2 (ΔH 0 - 297 kJ/mol);

SO 2 + ½ O 2 → SO 3 (ΔH 0 - 9.8 kJ/mol);

SO 3 + H 2 O → H 2 SO 4 (ΔH 0 - 130 kJ/mol).

Chemical properties of sulfuric acid

Sulfuric acid is a strong dibasic acid. In the first step, in solutions of low concentration, it dissociates almost completely:

H 2 SO 4 ↔H + + HSO 4 - .

Second stage dissociation

HSO 4 — ↔H + + SO 4 2-

occurs to a lesser extent. The dissociation constant of sulfuric acid in the second stage, expressed in terms of ion activity, K 2 = 10 -2.

As a dibasic acid, sulfuric acid forms two series of salts: medium and acidic. Average salts of sulfuric acid are called sulfates, and acid salts are called hydrosulfates.

Sulfuric acid greedily absorbs water vapor and is therefore often used to dry gases. The ability to absorb water also explains the charring of many organic matter, especially those belonging to the class of carbohydrates (fiber, sugar, etc.), when exposed to concentrated sulfuric acid. Sulfuric acid removes hydrogen and oxygen from carbohydrates, which form water, and carbon is released in the form of coal.

Concentrated sulfuric acid, especially hot, is a vigorous oxidizing agent. It oxidizes HI and HBr (but not HCl) to free halogens, coal to CO 2, sulfur to SO 2. These reactions are expressed by the equations:

8HI + H 2 SO 4 = 4I 2 + H 2 S + 4H 2 O;

2HBr + H 2 SO 4 = Br 2 + SO 2 + 2H 2 O;

C + 2H 2 SO 4 = CO 2 + 2SO 2 + 2H 2 O;

S + 2H 2 SO 4 = 3SO 2 + 2H 2 O.

The interaction of sulfuric acid with metals occurs differently depending on its concentration. Dilute sulfuric acid oxidizes with its hydrogen ion. Therefore, it interacts only with those metals that are in the voltage series only up to hydrogen, for example:

Zn + H 2 SO 4 = ZnSO 4 + H 2.

However, lead does not dissolve in dilute acid, since the resulting salt PbSO 4 is insoluble.

Concentrated sulfuric acid is an oxidizing agent due to sulfur (VI). It oxidizes metals in the voltage range up to and including silver. The products of its reduction may vary depending on the activity of the metal and the conditions (acid concentration, temperature). When interacting with low-active metals, such as copper, the acid is reduced to SO 2:

Cu + 2H 2 SO 4 = CuSO 4 + SO 2 + 2H 2 O.

When interacting with more active metals, reduction products can be both dioxide and free sulfur and hydrogen sulfide. For example, when interacting with zinc, the following reactions can occur:

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.

Application of sulfuric acid

The use of sulfuric acid varies from country to country and from decade to decade. For example, in the USA, the main area of ​​consumption of H 2 SO 4 is currently the production of fertilizers (70%), followed by chemical production, metallurgy, oil refining (~5% in each area). In the UK, the distribution of consumption by industry is different: only 30% of H2SO4 produced is used in the production of fertilizers, but 18% goes to paints, pigments and semi-products of dye production, 16% to chemical production, 12% to the production of soaps and detergents, 10 % for the production of natural and artificial fibers and 2.5% is used in metallurgy.

Examples of problem solving

EXAMPLE 1

Exercise	Determine the mass of sulfuric acid that can be obtained from one ton of pyrite if the yield of sulfur (IV) oxide in the roasting reaction is 90%, and sulfur (VI) oxide in the catalytic oxidation of sulfur (IV) is 95% of theoretical.
Solution	Let us write the equation for the pyrite firing reaction: 4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2. Let's calculate the amount of pyrite substance: n(FeS 2) = m(FeS 2) / M(FeS 2); M(FeS 2) = Ar(Fe) + 2×Ar(S) = 56 + 2×32 = 120g/mol; n(FeS 2) = 1000 kg / 120 = 8.33 kmol. Since in the reaction equation the coefficient for sulfur dioxide is twice as large as the coefficient for FeS 2, then the theoretically possible amount of sulfur oxide (IV) substance is equal to: n(SO 2) theor = 2 ×n(FeS 2) = 2 ×8.33 = 16.66 kmol. And the practically obtained amount of moles of sulfur (IV) oxide is: n(SO 2) pract = η × n(SO 2) theor = 0.9 × 16.66 = 15 kmol. Let us write the reaction equation for the oxidation of sulfur oxide (IV) to sulfur oxide (VI): 2SO 2 + O 2 = 2SO 3. The theoretically possible amount of sulfur oxide (VI) is equal to: n(SO 3) theor = n(SO 2) pract = 15 kmol. And the practically obtained amount of moles of sulfur oxide (VI) is: n(SO 3) pract = η × n(SO 3) theor = 0.5 × 15 = 14.25 kmol. Let us write the reaction equation for the production of sulfuric acid: SO 3 + H 2 O = H 2 SO 4. Let's find the amount of sulfuric acid: n(H 2 SO 4) = n(SO 3) pract = 14.25 kmol. The reaction yield is 100%. The mass of sulfuric acid is equal to: m(H 2 SO 4) = n(H 2 SO 4) × M(H 2 SO 4); M(H 2 SO 4) = 2×Ar(H) + Ar(S) + 4×Ar(O) = 2×1 + 32 + 4×16 = 98 g/mol; m(H 2 SO 4) = 14.25 × 98 = 1397 kg.
Answer	The mass of sulfuric acid is 1397 kg

Sulfur compounds (1U). Sulfurous acid

In tetrahalides SHal 4, oxohalides SOI Ial 2 and dioxide S0 2, sulfurous acid 1I 2 S0 3, sulfur exhibits an oxidation state of +4. In all these compounds, as well as in their corresponding anionic complexes, the sulfur atom has an unshared pair of electrons. Based on the number of a-bonding and non-bonding electrons, the shape of the molecules of these compounds changes from a distorted tetrahedron (SHal 4) to an angular shape (S0 9) through a trigonal pyramid shape (SOHal 2 and SO3). S(IV) compounds have acidic properties, which manifests themselves in reactions with water:

Sulfur oxide(1U) S0 2, or sulfur dioxide, is formed by burning sulfur in air or oxygen, as well as by calcining sulfides, such as pyrite:

Pyrite oxidation underlies the industrial method for producing S0 2. The S0 2 molecule is built similarly to the Oe molecule and has the structure isosceles triangle with a sulfur atom at the top. Length S-O connection is 0.143 nm, and the bond angle is 119.5°:

The sulfur atom is in the 5/? 2-hybridization. The p-orbital is oriented perpendicular to the plane of the molecule and is not involved in hybridization (Fig. 25.2). Due to this and other similarly oriented p-orbitals of oxygen atoms, a three-center n-bond is formed.

Rice. 25.2.

Under normal conditions, sulfur oxide (1U) is a colorless gas with a characteristic pungent odor. Let's dissolve well in water. Aqueous solutions have an acidic reaction, since S0 2, interacting with water, forms sulfurous acid H 2 S0 3. The reaction is reversible:

A characteristic feature of S0 2 is its redox duality. This is explained by the fact that in SO. ; sulfur has an oxidation state of +4, and therefore it can, by donating two electrons, be oxidized to S(VI), and by receiving four electrons, reduced to S. The manifestation of these and other properties depends on the nature of the reacting component. Thus, with strong oxidizing agents, S0 2 behaves as a typical reducing agent. For example, halogens are reduced to the corresponding hydrogen halides, and S(IV) usually transforms into S(VI):

In the presence of strong reducing agents, S0 2 behaves as an oxidizing agent:

It is also characterized by a disproportionation reaction:

SQ is an acidic oxide, easily soluble in water (1 volume of H 2 0 dissolves 40 volumes of S0 2). An aqueous solution of SOv is acidic and is called sulfurous acid. Typically, the bulk of S0 2 dissolved in water is in solution in the hydrated form of S0 2 azH 2 0, and only a small part of S0 2 interacts with water according to the scheme

Sulfurous acid, as a dibasic acid, forms two types of salts: medium - sulfites (Na 2 S0 3) and acidic - hydrosulfites (NaHS0 3). H 2 S0 3 exists in two tautomeric forms (Fig. 25.3).

Rice. 25.3.Structure of tautomeric forms of H 2 S0 3

Since sulfur in sulfurous acid has an oxidation state of +4, it exhibits, like S0 2, the properties of both an oxidizing agent and a reducing agent, as already mentioned, therefore sulfurous acid in oxidation-reduction reactions completely duplicates the properties of S0 9.

Salts H 2 S0 3 (sulfites) have the properties of both oxidizing and reducing agents. Thus, the SO 2 ion easily transforms into the SO 2 ion, exhibiting strong reducing properties, therefore, in solutions, sulfites are gradually oxidized by molecular oxygen, turning into sulfuric acid salts:

In the presence of strong reducing agents, sulfites behave as oxidizing agents. With strong heating, sulfites of the most active metals decompose at 600°C to form salts H 2 SO^ and H 2 S, i.e. disproportion occurs:

Of the salts of sulfurous acid, only salts of 5-elements of group I are dissolved, as well as hydrosulfites of the Me 2+ (HS0 3) 2 type.

Since H 2 S0 3 is a weak acid, when acids act on sulfites and hydrosulfites, S0 2 is released, which is usually used to obtain S0 2 in the laboratory:

Water-soluble sulfites easily undergo hydrolysis, as a result of which the concentration of OH ions in the solution increases:

When S0 2 is passed through aqueous solutions of hydrosulfites, pyrosulfites are formed:

If a solution of Na 2 S0 3 is boiled with sulfur powder, then sodium thiosulfate is formed. In thiosulfates, sulfur atoms are located in two different degrees oxidation - +6 and -2:

The resulting thiosulfate ion corresponds to the acid H 2 S 2 0 3, called thiosulfuric acid. The free acid is unstable under normal conditions and easily decomposes:

The properties of thiosulfates are due to the presence of and in them, and

the presence of S determines the reducing properties of the S 2 0 3 _ ion:

Weaker oxidizing agents oxidize sodium thiosulfate to tetrathionic acid salts. An example is the interaction with iodine:

This reaction is widely used in analytical chemistry, as it is the basis of one of the most important methods of volumetric analysis, called iodometry.

Thiosulfates alkali metals are produced industrially on a large scale. Among them highest value has sodium thiosulfate Na 2 S 2 0 3, which is used in medicine as an antidote for poisoning with halogens and cyanides. The action of this drug is based on its property of releasing sulfur, which, for example, with cyanide ions CN forms the less toxic thiocyanate ion SCN:

The drug can also be used for poisoning with compounds As, Pb, Hg, since non-toxic sulfides are formed. Na 2 S 2 0 3 is used for allergic diseases, arthritis, neuralgia. A characteristic reaction for Na 2 S 2 0 3 is its interaction with AgN0 3: a precipitate is formed white Ag. ; S.; 0 3, which over time under the influence of light and moisture turns black with the release of Ag 2 S:

These reactions are used for the qualitative detection of thiosulfate ion.

Thionyl chloride SOCl 2 is obtained by reacting S0 2 with PC1 5:

The SOCl 2 molecule has a pyramidal structure (Fig. 25.4). Bonds with sulfur are formed due to a set of orbitals, which can be very approximately considered as $/? 3-hybrid. One of them is occupied by a lone pair of electrons, so SOCl 2 can exhibit the properties of a weak Lewis base.

Rice. 25.4.

S()C1 2 - colorless, fuming liquid with a pungent odor, hydrolyzes in the presence of traces of moisture:

Volatile compounds formed during the reaction are easily removed. Therefore, SOCl 2 is often used to obtain anhydrous chlorides:

SOCl 2 is widely used as a chlorinating agent in organic syntheses.

Undiluted sulfuric acid is a covalent compound.

In the molecule, sulfuric acid is tetrahedrally surrounded by four oxygen atoms, two of which are part of the hydroxyl groups. The S–O bonds are double, and the S–OH bonds are single.

The colorless, ice-like crystals have a layered structure: each H 2 SO 4 molecule is connected to four neighboring strong hydrogen bonds, forming a single spatial framework.

The structure of liquid sulfuric acid is similar to the structure of solid one, only the integrity of the spatial framework is broken.

Physical properties of sulfuric acid

Under normal conditions, sulfuric acid is a heavy, oily liquid without color or odor. In technology, sulfuric acid is a mixture of both water and sulfuric anhydride. If the molar ratio of SO 3: H 2 O is less than 1, then it is an aqueous solution of sulfuric acid; if it is greater than 1, it is a solution of SO 3 in sulfuric acid.

100% H 2 SO 4 crystallizes at 10.45 ° C; T kip = 296.2 °C; density 1.98 g/cm3. H 2 SO 4 mixes with H 2 O and SO 3 in any ratio to form hydrates; the heat of hydration is so high that the mixture can boil, splash and cause burns. Therefore, it is necessary to add acid to water, and not vice versa, since when water is added to acid, lighter water will end up on the surface of the acid, where all the heat generated will be concentrated.

When aqueous solutions of sulfuric acid containing up to 70% H 2 SO 4 are heated and boiled, only water vapor is released into the vapor phase. Sulfuric acid vapor also appears above more concentrated solutions.

In terms of structural features and anomalies, liquid sulfuric acid is similar to water. There is the same system of hydrogen bonds, almost the same spatial framework.

Chemical properties of sulfuric acid

Sulfuric acid is one of the strongest mineral acids; due to its high polarity, the H–O bond is easily broken.

Sulfuric acid dissociates in aqueous solution , forming a hydrogen ion and an acidic residue:

H 2 SO 4 = H + + HSO 4 - ;

HSO 4 - = H + + SO 4 2- .

Summary equation:

H 2 SO 4 = 2H + + SO 4 2- .

Shows properties of acids , reacts with metals, metal oxides, bases and salts.

Dilute sulfuric acid does not exhibit oxidizing properties; when it interacts with metals, hydrogen and a salt containing the metal in the lowest oxidation state are released. In the cold, the acid is inert towards metals such as iron, aluminum and even barium.

Concentrated acid has oxidizing properties. Possible interaction products simple substances with concentrated sulfuric acid are given in the table. The dependence of the reduction product on the acid concentration and the degree of activity of the metal is shown: the more active the metal, the more deeply it reduces the sulfate ion of sulfuric acid.

Interaction with oxides:

CaO + H 2 SO 4 = CaSO 4 = H 2 O.

Interaction with bases:

2NaOH + H 2 SO 4 = Na 2 SO 4 + 2H 2 O.

Interaction with salts:

Na 2 CO 3 + H 2 SO 4 = Na 2 SO 4 + CO 2 + H 2 O.

Oxidative properties

Sulfuric acid oxidizes HI and HBr to free halogens:

H 2 SO 4 + 2HI = I 2 + 2H 2 O + SO 2.

Sulfuric acid takes away chemically bound water from organic compounds containing hydroxyl groups. Dehydration of ethyl alcohol in the presence of concentrated sulfuric acid leads to the production of ethylene:

C 2 H 5 OH = C 2 H 4 + H 2 O.

The charring of sugar, cellulose, starch and other carbohydrates upon contact with sulfuric acid is also explained by their dehydration:

C 6 H 12 O 6 + 12H 2 SO 4 = 18H 2 O + 12SO 2 + 6CO 2.

Sulfur dioxide is formed when sulfur is burned in air or oxygen. It is also obtained by calcining metal sulfides, such as iron pyrites, in air (“burning”):

By this reaction, sulfur dioxide is usually obtained in industry (about others industrial methods receipt see 9 § 131).

Sulfur dioxide is a colorless gas (“sulfur dioxide”) with a strong odor of hot sulfur. It condenses quite easily into a colorless liquid, boiling at . When a liquid evaporates, a strong decrease in temperature occurs (to ).

Sulfur dioxide is highly soluble in water (about 40 volumes in 1 volume of water at ); in this case, a partial reaction with water occurs and sulfurous acid is formed:

Thus, sulfur dioxide is an anhydride of sulfurous acid. When heated, solubility decreases and the equilibrium shifts to the left; gradually all the sulfur dioxide is released from the solution again.

The molecule is constructed similarly to the ozone molecule. The nuclei of its constituent atoms form an isosceles triangle:

Here the sulfur atom, like the central oxygen atom in the ozone molecule, is in a state of -hybridization and the angle is close to . The -orbital of the sulfur atom, oriented perpendicular to the plane of the molecule, does not participate in hybridization. Due to this orbital and similarly oriented -orbitals of oxygen atoms, a three-center -bond is formed; the pair of electrons that carry it out belongs to all three atoms of the molecule.

Sulfur dioxide is used to produce sulfuric acid, and also (in much smaller quantities) for bleaching straw, wool, silk and as a disinfectant (to destroy mold in basements, cellars, wine barrels, fermentation tanks).

Sulfurous acid is a very fragile compound. It is known only in aqueous solutions. When trying to separate sulfurous acid, it breaks down into water. For example, when concentrated sulfuric acid acts on sodium sulfite, sulfur dioxide is released instead of sulfurous acid:

The sulfurous acid solution must be protected from access to air, otherwise it, absorbing oxygen from the air, slowly oxidizes into sulfuric acid:

Sulfurous acid is a good reducing agent. For example, it reduces free halogens into hydrogen halides:

However, when interacting with strong reducing agents, sulfurous acid can play the role of an oxidizing agent. So, its reaction with hydrogen sulfide mainly proceeds according to the equation:

Being dibasic, sulfurous acid forms two series of salts. Its average salts are called sulfites, acidic ones - hydrosulfites.

Like acid, sulfites and hydrosulfites are reducing agents. When they are oxidized, salts of sulfuric acid are obtained.

When calcined, sulfites of the most active metals decompose to form sulfides and sulfates (self-oxidation - self-healing reaction):

Potassium and sodium sulfites are used for bleaching certain materials, in the textile industry for dyeing fabrics, and in photography. The solution (this salt exists only in solution) is used to process wood into so-called sulfite pulp, from which paper is then obtained.