Part I

1. Hydrogen sulfide.
1) Molecule structure:

2) Physical properties: a colorless gas with a pungent odor of rotten eggs, heavier than air.

3) Chemical properties (complete the reaction equations and consider the equations in the light of TED or from the standpoint of oxidation-reduction).

4) Hydrogen sulfide in nature: in the form of compounds - sulfides, in free form - in volcanic gases.

2. Sulfur (IV) oxide – SO2
1) Obtained in industry. Write down the reaction equations and consider them from the standpoint of oxidation-reduction.

2) Obtained in the laboratory. Write down the reaction equation and consider it in the light of TED:

3) Physical properties: gas with a pungent suffocating odor.

4) Chemical properties.

3. Sulfur oxide (VI) - SO3.
1) Preparation by synthesis from sulfur oxide (IV):

2) Physical properties: liquid, heavier than water, mixed with sulfuric acid - oleum.

3) Chemical properties. Exhibits typical properties of acid oxides:

Part II

1. Characterize the reaction for the synthesis of sulfur oxide (VI) according to all classification criteria.

a) catalytic
b) reversible
c) OVR
d) connections
e) exothermic
e) combustion

2. Characterize the reaction of sulfur (IV) oxide with water according to all classification criteria.

a) reversible
b) connections
c) not OVR
d) exothermic
e) non-catalytic

3. Explain why hydrogen sulfide exhibits strong reducing properties.

4. Explain why sulfur (IV) oxide can exhibit both oxidizing and reducing properties:

Confirm this thesis with the equations of the corresponding reactions.

5. Sulfur of volcanic origin is formed as a result of the interaction of sulfur dioxide and hydrogen sulfide. Write down the reaction equations and consider them from the standpoint of oxidation-reduction.

6. Write down the equations of transition reactions, deciphering the unknown formulas:

7. Write a syncwine on the topic “Sulfur dioxide.”
1) Sulfur dioxide
2) Suffocating and harsh
3) Acid oxide, OVR
4) Used to produce SO3
5) Sulfuric acid H2SO4

8. Using additional sources of information, including the Internet, prepare a message about the toxicity of hydrogen sulfide (pay attention to its characteristic smell!) and first aid for poisoning with this gas. Write down your message plan in a special notebook.

Hydrogen sulfide
Colorless gas with the smell of rotten eggs. It is detected in the air by smell even in small concentrations. In nature, it is found in water from mineral springs, seas, and volcanic gases. Formed during the decomposition of proteins without access to oxygen. It can be released into the air in a number of chemical and textile industries, during oil production and refining, and from sewage systems.
Hydrogen sulfide is a strong poison that causes acute and chronic poisoning. It has a local irritant and general toxic effect. At a concentration of 1.2 mg/l, poisoning develops at lightning speed, death occurs due to acute inhibition of tissue respiration processes. When exposure is stopped, even in severe forms of poisoning, the victim can be brought back to life.
At a concentration of 0.02-0.2 mg/l, headache, dizziness, chest tightness, nausea, vomiting, diarrhea, loss of consciousness, convulsions, damage to the mucous membrane of the eyes, conjunctivitis, photophobia are observed. The risk of poisoning increases due to loss of smell. Cardiac weakness and respiratory failure, coma gradually increase.
First aid - removing the victim from the polluted atmosphere, inhaling oxygen, artificial respiration; means that stimulate the respiratory center, warming the body. Glucose, vitamins, and iron supplements are also recommended.
Prevention - sufficient ventilation, sealing of some production operations. When lowering workers into wells and containers containing hydrogen sulfide, they must use gas masks and life belts on ropes. Gas rescue service is mandatory in mines, production sites and oil processing plants.

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 the reducing agent SO 2 reacts 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 is widely used in the textile industry for bleaching various products. In addition, it is used in agriculture to destroy harmful microorganisms in greenhouses and cellars. Large quantities of SO 2 are 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 its pure form, sulfuric anhydride has no practical significance. It is obtained as an intermediate product in the production of sulfuric acid.

H2SO4

Mention of sulfuric acid is 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 Libavius ​​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 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 exhibits all the characteristic properties of 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.

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 with a diagram:

H2SO4∙ n SO 3 + H 2 O = H 2 SO 4

Sulfuric acid finds a variety of applications in a wide variety of areas of the 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

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 + 2H2O

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 aqueous solution, is formed:

SO 2 +H 2 O «H 2 SO 3

The reaction is reversible. In an aqueous solution, sulfur oxide (IV) and sulfurous acid are in chemical equilibrium, which can be displaced. When binding H 2 SO 3 (neutralization of acid

u) 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 materials. 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 when producing 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 + NON «NaHSO 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, resulting in the color being restored. Therefore, 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. It is also used for the 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.

The +4 oxidation state for sulfur is quite stable and manifests itself in SHal 4 tetrahalides, SOHal 2 oxodihalides, SO 2 dioxide and their corresponding anions. We will get acquainted with the properties of sulfur dioxide and sulfurous acid.

1.11.1. Sulfur (IV) oxide Structure of the so2 molecule

The structure of the SO 2 molecule is similar to the structure of the ozone molecule. The sulfur atom is in a state of sp 2 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.

Physical properties

Under normal conditions, sulfur dioxide or sulfur dioxide is a colorless gas with a sharp suffocating odor, melting point -75 °C, boiling point -10 °C. It is highly soluble in water; at 20 °C, 40 volumes of sulfur dioxide dissolve in 1 volume of water. Toxic gas.

Chemical properties of sulfur (IV) oxide

Sulfur dioxide is highly reactive.

Sulfur dioxide is an acidic oxide. It is quite soluble in water to form hydrates. It also partially reacts with water, forming weak sulfurous acid, which is not isolated in individual form:

SO 2 + H 2 O = H 2 SO 3 = H + + HSO 3 - = 2H + + SO 3 2- .

As a result of dissociation, protons are formed, so the solution has an acidic environment.

When sulfur dioxide gas is passed through a sodium hydroxide solution, sodium sulfite is formed. Sodium sulfite reacts with excess sulfur dioxide to form sodium hydrosulfite:

2NaOH + SO 2 = Na 2 SO 3 + H 2 O;

Na 2 SO 3 + SO 2 = 2NaHSO 3.

Sulfur dioxide is characterized by redox duality; for example, it exhibits reducing properties and decolorizes bromine water:

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

and potassium permanganate solution:

5SO 2 + 2KMnO 4 + 2H 2 O = 2KНSO 4 + 2MnSO 4 + H 2 SO 4.

oxidized by oxygen to sulfuric anhydride:

2SO 2 + O 2 = 2SO 3.

It exhibits oxidizing properties when interacting with strong reducing agents, for example:

SO 2 + 2CO = S + 2CO 2 (at 500 °C, in the presence of Al 2 O 3);

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

Preparation of sulfur oxide (IV)

Combustion of sulfur in air

S + O 2 = SO 2.

Sulfide oxidation

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

Effect of strong acids on metal sulfites

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

1.11.2. Sulfurous acid and its salts

When sulfur dioxide is dissolved in water, weak sulfurous acid is formed, the bulk of dissolved SO 2 is in the form of the hydrated form SO 2 ·H 2 O; upon cooling, crystalline hydrate is also released, only a small part of the sulfurous acid molecules dissociates into sulfite and hydrosulfite ions. In the free state, the acid is not released.

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 fairly 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/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 with 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 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 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 valence. May be highly toxic and cause serious health problems.