Pyrite sulfur oxide 4. Sulfur dioxide - physical properties, production and application

Sulfur oxide (IV) has acidic properties, which are manifested in reactions with substances that exhibit basic properties. Acidic properties are manifested when interacting with water. In this case, a solution of sulfuric acid is formed:

The oxidation state of sulfur in sulfur dioxide (+4) determines the reducing and oxidizing properties of sulfur dioxide:

vo-tel: S + 4 - 2e => S + 6

oct: S+4 + 4e => S0

Reducing properties are manifested in reactions with strong oxidizing agents: oxygen, halogens, nitric acid, potassium permanganate and others. For example:

2SO2 + O2 = 2SO3

S+4 - 2e => S+6 2

O20 + 4e => 2O-2 1

With strong reducing agents, the gas exhibits oxidizing properties. For example, if you mix sulfur dioxide and hydrogen sulfide, they interact under normal conditions:

2H2S + SO2 = 3S + 2H2O

S-2 - 2e => S0 2

S+4 + 4e => S0 1

Sulfurous acid exists only in solution. It is unstable and decomposes into sulfur dioxide and water. Sulfurous acid is not a strong acid. It is an acid of medium strength and dissociates in steps. When alkali is added to sulfuric acid, salts are formed. Sulfurous acid gives two series of salts: medium - sulfites and acidic - hydrosulfites.

Sulfur(VI) oxide

Sulfur trioxide exhibits acidic properties. It reacts violently with water, and a large amount of heat is released. This reaction is used to obtain the most important product chemical industry- sulfuric acid.

SO3 + H2O = H2SO4

Since sulfur in sulfur trioxide has the highest oxidation state, sulfur(VI) oxide exhibits oxidizing properties. For example, it oxidizes halides, non-metals with low electronegativity:

2SO3 + C = 2SO2 + CO2

S+6 + 2e => S+4 2

C0 - 4e => C+4 2

Sulfuric acid reacts three types: acid-base, ion-exchange, redox. It also actively interacts with organic substances.

Acid-base reactions

Sulfuric acid exhibits acidic properties in reactions with bases and basic oxides. These reactions are best carried out with dilute sulfuric acid. Insofar as sulphuric acid is dibasic, it can form both medium salts (sulfates) and acidic salts (hydrosulfates).

Ion exchange reactions

Sulfuric acid is characterized by ion exchange reactions. At the same time, it interacts with salt solutions, forming a precipitate, a weak acid, or releasing a gas. These reactions proceed at a faster rate when using 45% or even more dilute sulfuric acid. Gas evolution occurs in reactions with salts of unstable acids, which decompose to form gases (carbonic, sulfurous, hydrogen sulfide) or to form volatile acids, such as hydrochloric.

Redox reactions

Sulfuric acid most clearly manifests its properties in redox reactions, since sulfur in its composition has the highest oxidation state of +6. The oxidizing properties of sulfuric acid can be found in the reaction, for example, with copper.

There are two oxidizing elements in a sulfuric acid molecule: a sulfur atom with S.O. +6 and hydrogen ions H+. Copper cannot be oxidized by hydrogen to the +1 oxidation state, but sulfur can. This is the reason for the oxidation of such an inactive metal as copper with sulfuric acid.

Sulfur dioxide has a molecular structure similar to ozone. The sulfur atom in the center of the molecule is bonded to two oxygen atoms. This gaseous product of sulfur oxidation is colorless, emits a pungent odor, easily condenses into a clear liquid under changing conditions. The substance is highly soluble in water, has antiseptic properties. IN large quantities receive SO 2 in the chemical industry, namely in the cycle of sulfuric acid production. 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. Official name connections, chemical composition which reflects the formula SO 2 - sulfur dioxide. IUPAC recommends the use of this term and its English equivalent, Sulfur dioxide. Textbooks for schools and universities often mention another name - sulfur oxide (IV). The Roman numeral in brackets denotes the valency of the S atom. The oxygen in this oxide is bivalent, and the oxidation number of sulfur is +4. The technical literature uses such obsolete terms as sulfur dioxide, sulfurous anhydride (the product of its dehydration).

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 involved in education. covalent bond between sulfur and oxygen. In the O–S pair, the distance between the atoms is 0.143 nm. Oxygen is more electronegative than sulfur, which means that the bonding pairs of electrons move from the center to the outer corners. The whole molecule is also polarized, the negative pole is the O atoms, the positive one is the S atom.

Some physical parameters of sulfur dioxide

Quadrivalent sulfur oxide at normal rates environment retains a gaseous state of aggregation. The sulfur dioxide formula allows you to determine its relative molecular and molar mass: Mr(SO 2) \u003d 64.066, M \u003d 64.066 g / mol (can be rounded up 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, causes a cough. But sulfur oxide (IV) is not as toxic as hydrogen sulfide.

under pressure at room temperature gaseous sulfur dioxide is liquefied. At low temperatures the substance is in a solid state, melts at -72 ... -75.5 ° C. With a further increase in temperature, a liquid appears, and at -10.1 ° C, gas is again formed. SO 2 molecules are thermally stable, decomposition into atomic sulfur and molecular oxygen occurs at very high temperatures (about 2800 ºС).

Solubility and interaction with water

Sulfur dioxide, when dissolved in water, partially interacts with it to form a very weak sulfurous acid. At the time 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. Gaseous dioxide interacts better with cool water, its solubility decreases with increasing temperature. Under normal conditions, it can dissolve in 1 volume of water up to 40 volumes of gas.

Sulfur dioxide in nature

Significant volumes of sulfur dioxide are released with volcanic gases and lava during eruptions. Many human activities also increase the concentration of SO 2 in the atmosphere.

Sulfur dioxide is supplied to the air by metallurgical plants, where exhaust gases are not captured during the roasting of ore. Many fossil fuels contain sulfur, resulting in the release of significant amounts of sulfur dioxide into atmospheric air when burning coal, oil, gas, fuel obtained from them. Sulfur dioxide becomes toxic to humans at concentrations in the air above 0.03%. A person begins shortness of breath, there may be phenomena resembling bronchitis and pneumonia. A very high concentration 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 \u003d SO 2.
You can act on the salts of sulfurous acid with stronger inorganic acids, it is better to take hydrochloric, but you can dilute sulfuric:

Na 2 SO 3 + 2HCl \u003d 2NaCl + H 2 SO 3;
Na 2 SO 3 + H 2 SO 4 (diff.) \u003d Na 2 SO 4 + H 2 SO 3;
H 2 SO 3 \u003d H 2 O + SO 2.

3. When copper interacts with concentrated sulfuric acid, not hydrogen is released, but sulfur dioxide:

2H 2 SO 4 (conc.) + Cu \u003d CuSO 4 + 2H 2 O + SO 2.

Modern ways industrial production sulfur dioxide:

Oxidation of natural sulfur during its combustion in special furnaces: S + O 2 = SO 2.
Roasting 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 interacts with sulfur dioxide, the reaction products are sulfuric acid and hydrogen bromide. The oxidizing properties of SO 2 are manifested if this gas is passed through hydrogen sulfide water. As a result, sulfur is released, self-oxidation-self-healing occurs: SO 2 + 2H 2 S \u003d 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; it is possible to detect the acidic properties of a sulfur dioxide solution using indicators (litmus turns pink). Sulfurous acid gives medium salts - sulfites and acidic - hydrosulfites. Among them are stable compounds.

The process of oxidation of sulfur in dioxide to a hexavalent state in sulfuric anhydride is catalytic. The resulting substance dissolves vigorously in water, reacts with H 2 O molecules. The reaction is exothermic, sulfuric acid is formed, or rather, its hydrated form.

Practical use of sour gas

The main process for the industrial production of sulfuric acid, which requires element 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 for the production of 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, the main starting material for sulfuric acid production in last years became natural sulfur. Significant world reserves of this raw material, 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 interaction to bleach silk and woolen fabrics. This is one of the types of chlorine-free bleaching, in which the fibers are not destroyed.

Sulfur dioxide has excellent disinfectant properties, which is used in the fight against fungi and bacteria. Sulfur dioxide is used to fumigate agricultural storages, wine barrels and cellars. Used by SO 2 in Food Industry as a preservative and antibacterial substance. Add it to syrups, soak fresh fruits in it. Sulfitization
sugar beet juice discolors and disinfects raw materials. Canned vegetable puree and juices also contain sulfur dioxide as an antioxidant and preservative agent.

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 oxide (IV) The 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, the shape of the molecule is angular. The sulfur atom has an unshared electron pair. The S-O bond length is 0.143 nm, the bond angle is 119.5°.

The structure corresponds to the following resonant structures:

Unlike ozone, the S–O bond multiplicity is 2, i.e., the first resonance structure makes the main contribution. The molecule is characterized by high thermal stability.

Physical properties

Under normal conditions, sulfur dioxide or sulfur dioxide is colorless gas with a sharp suffocating odor, melting point -75 °C, boiling point -10 °C. Let's well dissolve in water, at 20 °C in 1 volume of water 40 volumes of sulfur dioxide are dissolved. Toxic gas.

Chemical properties of sulfur oxide (IV)

Sulfur dioxide is highly reactive. Sulfur dioxide is an acid oxide. It is quite soluble in water with the formation of hydrates. It also partially interacts with water, forming a weak sulfurous acid, which is not isolated individually:

SO 2 + H 2 O \u003d H 2 SO 3 \u003d H + + HSO 3 - \u003d 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 \u003d 2NaHSO 3.

Sulfur dioxide is characterized by redox duality, for example, it, showing reducing properties, discolors bromine water:

SO 2 + Br 2 + 2H 2 O \u003d H 2 SO 4 + 2HBr

and potassium permanganate solution:

5SO 2 + 2KMnO 4 + 2H 2 O \u003d 2KНSO 4 + 2MnSO 4 + H 2 SO 4.

oxidized by oxygen to sulfuric anhydride:

2SO 2 + O 2 \u003d 2SO 3.

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

SO 2 + 2CO \u003d S + 2CO 2 (at 500 ° C, in the presence of Al 2 O 3);

SO 2 + 2H 2 \u003d S + 2H 2 O.

Production of sulfur oxide (IV)

Burning sulfur in air

S + O 2 \u003d SO 2.

Sulfide oxidation

4FeS 2 + 11O 2 \u003d 2Fe 2 O 3 + 8SO 2.

The action of strong acids on metal sulfites

Na 2 SO 3 + 2H 2 SO 4 \u003d 2NaHSO 4 + H 2 O + SO 2.

1.11.2. Sulfuric acid and its salts

When sulfur dioxide is dissolved in water, weak sulfurous acid is formed, the bulk of the dissolved SO 2 is in the form of a hydrated form of SO 2 H 2 O, upon cooling, a 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 isolated.

Being dibasic, it forms two types of salts: medium - sulfites and acidic - hydrosulfites. Only alkali metal sulfites and hydrosulfites of alkali and alkaline earth metals dissolve in water.

4.doc

Sulfur. Hydrogen sulfide, sulfides, hydrosulfides. Sulfur (IV) and (VI) oxides. 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 widely distributed in nature and occurs both in the free state (native sulfur) and in compounds. Minerals containing sulfur can be divided into two groups - sulfides (pyrites, shines, blendes) 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 the Crimea and other regions.

The 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.

The minerals of the second group include gypsum CaSO 4 2H 2 O, mirabilite (Glauber's salt) - Na 2 SO 4 10H 2 O, ki-serite - MgSO 4 H 2 O.

Sulfur is found in organisms of animals and plants, as it is part of protein molecules. Organic sulfur compounds are found in oil.

Receipt

1. When receiving sulfur from natural compounds, for example, from sulfur pyrites, it is heated to high temperatures. Sulfur pyrite decomposes with the formation of iron (II) sulfide and sulfur:

2. Sulfur can be obtained by the oxidation of hydrogen sulfide with a lack of oxygen according to the reaction:

2H 2 S + O 2 \u003d 2S + 2H 2 O

3. Currently, it is common to obtain sulfur by carbon reduction of sulfur dioxide SO 2 - a by-product in the smelting of metals from sulfur ores:

SO 2 + C \u003d CO 2 + S

4. Off-gases from metallurgical and coke ovens contain a mixture of sulfur dioxide and hydrogen sulfide. This mixture is passed at high temperature over a catalyst:

H 2 S + SO 2 \u003d 2H 2 O + 3S

^ Physical properties

Sulfur is a brittle solid lemon yellow. It is practically insoluble in water, but highly soluble in carbon disulfide CS 2 aniline and some other solvents.

Poor conductor of heat and electricity. Sulfur forms several allotropic modifications:

1 . ^ Rhombic sulfur (the most stable), crystals have the form of octahedrons.

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 formed when 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 unfinished outer 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 giving or pulling electrons 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. During these reactions, sulfides are formed (does not react with gold, platinum and iridium): Fe + S = FeS

2. Under normal conditions, sulfur does not interact with hydrogen, and at 150-200 ° C a reversible reaction occurs:

3. In reactions with metals and hydrogen, sulfur behaves like a typical oxidizing agent, and in the presence of strong oxidizing agents it exhibits reducing properties.

S + 3F 2 \u003d SF 6 (does not react with iodine)

4. The combustion of sulfur in oxygen proceeds at 280°C, and in air at 360°C. This forms a mixture of SO 2 and SO 3:

S + O 2 \u003d SO 2 2S + 3O 2 \u003d 2SO 3

5. When heated without air access, sulfur directly combines with phosphorus, carbon, showing oxidizing properties:

2P + 3S \u003d P 2 S 3 2S + C \u003d CS 2

6. When interacting with complex substances sulfur behaves primarily as a reducing agent:

7. Sulfur is capable of disproportionation reactions. So, 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, which turns the rubber into rubber.

In the form of a sulfur color (fine powder), sulfur is used to combat diseases of the vineyard and cotton. It is used to obtain gunpowder, matches, luminous compositions. In medicine, sulfur ointments are prepared for the treatment of skin diseases.

4.2. Hydrogen sulfide, sulfides, hydrosulfides

Hydrogen sulfide is analogous to water. Its electronic formula

Shows that in education H-S-H bonds two p-electrons involved external level sulfur atom. 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, such as 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 usually it is obtained by the action of dilute hydrochloric or sulfuric acid on iron (III) sulfide:

2HCl+FeS=FeCl 2 +H 2 S 2H + +FeS=Fe 2+ +H 2 S This reaction is often carried out in a Kipp apparatus.

^ Physical properties

Under normal conditions, hydrogen sulfide is a colorless gas with a strong characteristic smell of rotten eggs. Very toxic, when inhaled, it binds to hemoglobin, causing paralysis, which is not uncommon.

Ko leads to death. Less dangerous in low concentrations. You need to work with him fume hoods or with hermetically sealed devices. Permissible content of H 2 S in industrial premises is 0.01 mg per 1 liter of air.

Hydrogen sulfide is relatively well 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, With strong heating, hydrogen sulfide almost completely decomposes with the formation of sulfur and hydrogen.

2. Gaseous hydrogen sulfide burns in air with a blue flame to form sulfur oxide (IV) and water:

2H 2 S + 3O 2 \u003d 2SO 2 + 2H 2 O

With a lack of oxygen, sulfur and water are formed: 2H 2 S + O 2 \u003d 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 a solution of H 2 S, it is relatively easy to donate electrons to air oxygen molecules:

At the same time, air oxygen oxidizes hydrogen sulfide to sulfur, which makes hydrogen sulfide water cloudy:

2H 2 S + O 2 \u003d 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 - atmospheric oxygen 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: under the action nitric acid free sulfur is formed.

6. A solution of hydrogen sulfide has an acidic reaction due to dissociations:

H 2 SH + +HS - HS - H + +S -2

Usually the first stage dominates. It is a very weak acid: weaker than carbonic, which usually displaces H 2 S from sulfides.

Sulfides and hydrosulfides

Hydrosulfuric acid, as dibasic, forms two series of salts:

Medium - sulfides (Na 2 S);

Acidic - hydrosulfides (NaHS).

These salts can be obtained: - by the interaction of hydroxides with hydrogen sulfide: 2NaOH + H 2 S = Na 2 S + 2H 2 O

By direct interaction of sulfur with metals:

Exchange reaction of salts with H 2 S or between salts:

Pb(NO 3) 2 + Na 2 S \u003d PbS + 2NaNO 3

CuSO 4 +H 2 S=CuS+H 2 SO 4 Cu 2+ +H 2 S=CuS+2H +

Almost all hydrosulfides are highly soluble in water.

Sulfides of alkali and alkaline earth metals are also easily soluble in water, 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, sulfides alkali metals when dissolved in water, they have an alkaline reaction:

Na 2 S+HOHNaHS+NaOH

All sulfides, like hydrogen sulfide itself, are energetic reducing agents:

3PbS -2 + 8HN +5 O 3 (razb.) \u003d 3PbS +6 O 4 + 4H 2 O + 8N +2 O

Some sulfides have a characteristic color: CuS and PbS - black, CdS - yellow, ZnS - white, MnS - pink, SnS - brown, Al 2 S 3 - orange. The qualitative analysis of cations is based on the different solubility of sulfides and the different colors of many of them.

^ 4.3. Sulfur(IV) oxide and sulfurous acid

Sulfur oxide (IV), or sulfur dioxide, under normal conditions, a colorless gas with a pungent suffocating odor. When cooled to -10°C, it liquefies into a colorless liquid.

Receipt

1. Under laboratory conditions, sulfur oxide (IV) is obtained from salts of sulfurous acid by the action of strong acids on them:

Na 2 SO 3 + H 2 SO 4 \u003d Na 2 SO 4 + S0 2  + H 2 O 2NaHSO 3 + H 2 SO 4 \u003d Na 2 SO 4 + 2SO 2  + 2H 2 O 2HSO - 3 + 2H + \u003d 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 \u003d CuSO 4 + SO 2  + 2H 2 O

Cu + 4Н + + 2SO 2- 4 \u003d Cu 2+ + SO 2- 4 + SO 2  + 2H 2 O

3. Sulfur oxide (IV) is also formed when sulfur is burned in air or oxygen:

4. Under industrial conditions, SO 2 is obtained by roasting pyrite FeS 2 or sulfurous ores of non-ferrous metals (zinc blende ZnS, lead luster PbS, etc.):

4FeS 2 + 11O 2 \u003d 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 the SO 2 molecule. The mutual repulsion of the bonding electron pairs and the non-shared electron pair of sulfur gives the molecule an angular shape.

Chemical properties

1. Sulfur oxide (IV) exhibits all the properties of acidic oxides:

Interaction with water

Interaction with alkalis,

Interaction with basic oxides.

2. Sulfur oxide (IV) is characterized by reducing properties:

S +4 O 2 +O 0 2 2S +6 O -2 3 (in the presence of a catalyst, when heated)

But in the presence of strong reducing agents, SO 2 behaves like 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, giving 2 electrons, be oxidized to S +6, and receiving 4 electrons, be 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 are dissolved in 1 volume at 20 ° C). In this case, sulfurous acid exists only in an aqueous solution:

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 moved. When binding H 2 SO 3 (neutralization of acid

You) the reaction proceeds towards the formation of sulfurous acid; when removing SO 2 (blowing through a nitrogen solution or heating), the reaction proceeds towards the starting materials. In a solution of sulfurous acid, there is always sulfur oxide (IV), which gives it a pungent odor.

Sulfurous acid has all the properties of acids. In the solution, it dissociates in steps:

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 \u003d Na 2 SO 3 + 2H 2 O

Hydrosulfites are obtained with a lack of alkali:

H 2 SO 3 + NaOH \u003d 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. So, under the action of oxygen, sulfites are oxidized to sulfates:

2Na 2 S +4 O 3 + O 0 2 \u003d 2Na 2 S +6 O -2 4

The oxidation of sulfurous acid with bromine and potassium permanganate proceeds even more easily:

5H 2 S +4 O 3 +2KMn +7 O 4 \u003d 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:

Salts of sulfurous acid dissolve almost all hydro-sulfites and sulfites of alkali metals.

3. Since H 2 SO 3 is a weak acid, the action of acids on sulfites and hydrosulfites releases SO 2. This method is usually used when obtaining SO 2 in laboratory conditions:

NaHSO 3 + H 2 SO 4 \u003d 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 increases in the solution:

Na 2 SO 3 + NOHNaHSO 3 + NaOH

Application

Sulfur oxide (IV) and sulfurous acid decolorize many dyes, forming colorless compounds with them. The latter can decompose again when heated or in the light, as a result of which the color is restored. Therefore, the whitening action of SO 2 and H 2 SO 3 differs from the whitening action of chlorine. Usually, sulfur (IV) rxide whitens wool, silk and straw.

Sulfur oxide (IV) kills many microorganisms. Therefore, to destroy mold fungi, they fumigate damp cellars, cellars, wine barrels, etc. It is also used in the transportation and storage of fruits and berries. In large quantities, sulfur oxide IV) is used to produce sulfuric acid.

Important application finds a solution of calcium hydrosulfite CaHSO 3 (sulfite liquor), which is used to treat wood and paper pulp.

^ 4.4. Sulfur(VI) oxide. Sulphuric 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 \u003d H 2 SO 4

Table 20. Properties of sulfur oxides

The dissolution of sulfur oxides (VI) in water is accompanied by the release of a significant amount of heat.

Sulfur oxide (VI) is very soluble in concentrated sulfuric acid. A solution of SO3 in anhydrous acid is called oleum. Oleums can contain up to 70% SO 3 .

Receipt

1. Sulfur oxide (VI) is produced by the oxidation of sulfur dioxide with atmospheric oxygen in the presence of catalysts at a temperature of 450 ° C (see. Getting sulfuric acid):

2SO 2 +O 2 \u003d 2SO 3

2. Another way to oxidize SO 2 to SO 3 is to use nitric oxide (IV) as an oxidizing agent:

The resulting nitric oxide (II) when interacting with atmospheric oxygen easily and quickly turns into nitric oxide (IV): 2NO + O 2 \u003d 2NO 2

Which again can be used in the oxidation of SO 2 . Therefore, NO 2 acts as an oxygen carrier. This method of oxidizing SO 2 to SO 3 is called nitrous. The 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 the binding electron pairs. The sulfur atom provided six external electrons for their formation.

Chemical properties

1. SO 3 is a typical acidic oxide.

2. Sulfur oxide (VI) has the properties of a strong oxidizing agent.

Application

Sulfur oxide (VI) is used to produce sulfuric acid. Highest value It has contact method receiving

Sulfuric acid. By this method, you can get H 2 SO 4 of any concentration, as well as oleum. The process consists of three stages: getting SO 2 ; oxidation of SO 2 to SO 3; getting H 2 SO 4 .

SO 2 is obtained by firing pyrite FeS 2 in special furnaces: 4FeS 2 + 11O 2 \u003d 2Fe 2 O 3 + 8SO 2

To speed up the firing, pyrite is preliminarily crushed, and for a more complete burnout of sulfur, much more air (oxygen) is introduced than is required by the reaction. The gas leaving the kiln consists of sulfur oxide (IV), oxygen, nitrogen, arsenic compounds (from impurities in pyrites) and water vapor. It is called roasting gas.

The roasting gas is thoroughly cleaned, 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 electro-filters and a washing tower; moisture is absorbed by concentrated sulfuric acid in the drying tower. The purified gas containing oxygen is heated in a heat exchanger up to 450°C and enters the contact apparatus. Inside the contact apparatus there are lattice shelves filled with a catalyst.

Previously, finely divided metallic 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

Increasing the oxygen content in the roasting gas increases the yield of sulfur oxide (VI): at a temperature of 450°C, it usually reaches 95% or more.

The resulting sulfur oxide (VI) is then fed countercurrently into the absorption tower, where it is absorbed by concentrated sulfuric acid. As it saturates, 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 ( \u003d 1.83 g / cm 3) is heavier than water, odorless, non-volatile. Extremely gigroscopic. Absorbs moisture with the release of a large amount of heat, so you can not add water to concentrated sulfuric acid - acid will splash. For times-

Additions of sulfuric acid should be added in small portions to water.

Anhydrous sulfuric acid dissolves up to 70% sulfur oxide (VI). When heated, it splits off SO 3 until a solution is formed with a mass fraction of H 2 SO 4 98.3%. Anhydrous H 2 SO 4 almost does not conduct electricity.

^ Chemical properties

1. It mixes with water in any ratio and forms hydrates of various composition:

H 2 SO 4 H 2 O, H 2 SO 4 2H 2 O, H 2 SO 4 3H 2 O, H 2 SO 4 4H 2 O, H 2 SO 4 6.5H 2 O

2. Concentrated sulfuric acid carbonizes organic substances - sugar, paper, wood, fiber, taking water elements from them:

C 12 H 22 O 11 + H 2 SO 4 \u003d 12C + H 2 SO 4 11H 2 O

The resulting coal partially interacts with the acid:

The drying of gases 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 \u003d NaHSO 4 + HNO 3

However, if you add H 2 SO 4 to salt solutions, then the displacement of acids does not occur.

H 2 SO 4 - strong dibasic acid: H 2 SO 4 H + + HSO - 4 HSO - 4 H + + SO 2- 4

It has all the properties of non-volatile strong acids.

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 voltages of metals up to hydrogen:

Interaction with metals is due to the reduction of hydrogen ions.

6. Concentrated sulfuric acid is a vigorous oxidizing agent. When heated, it oxidizes most metals, including those standing in the electrochemical series of voltages after hydrogen. It does not react only with platinum and gold. Depending on the activity of the metal, S -2 , S° and S +4 can be used as reduction products.

In the cold, concentrated sulfuric acid does not interact with such strong metals as aluminum, iron, chromium. This is due to the passivation of metals. This feature is widely used when transporting it in an iron container.

However, when heated:

Thus, concentrated sulfuric acid interacts with metals by reducing the atoms of the acid-forming agent.

A qualitative reaction to the sulfate ion SO 2- 4 is the formation of a white crystalline precipitate BaSO 4, insoluble in water and acids:

SO 2- 4 + Ba +2 BaSO 4 

Application

Sulfuric acid is the most important product the main chemical industry engaged in the production of non-

organic acids, alkalis, salts, mineral fertilizers and chlorine.

In terms of the variety of applications, sulfuric acid occupies the first place among acids. The largest number it is used to obtain phosphorus and nitrogen fertilizers. Being non-volatile, sulfuric acid is used to obtain other acids - hydrochloric, hydrofluoric, phosphoric and acetic.

A lot of it goes to the purification of 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, chromium 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 - potash 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 doughy mass, after a while hardens again, turning into CaSO 4 2H 2 O. Gypsum is widely used in construction (plaster).

^ Magnesium sulfate MgSO 4 is found in sea water, causing its bitter taste. The crystalline hydrate, called bitter salt, is used as a laxative.

vitriol- the technical name of crystalline sulphates of metals Fe, Cu, Zn, Ni, Co (dehydrated salts are not vitriol). blue vitriol CuSO 4 5H 2 O - poisonous substance of blue color. Plants are sprayed with a diluted solution and seeds are dressed before sowing. inkstone FeSO 4 7H 2 O is a light green substance. Used for plant pest control, preparation of inks, mineral paints, etc. Zinc vitriol ZnSO 4 7H 2 O is used in the production of mineral paints, in chintz printing, and medicine.

^ 4.5. Esters of sulfuric acid. Sodium thiosulfate

Sulfuric acid esters include dialkyl sulfates (RO 2)SO 2 . These are high-boiling liquids; the lower ones are soluble in water; in the presence of alkalis, they form alcohol and salts of sulfuric acid. 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 the interaction of sulfuric acid with ethanol. It is an ethylating 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. Let's dissolve in alcohols, it is bad - in water. Reacts with ammonia in the absence of a solvent (explosively); sulfonates some aromatic compounds, such as phenol esters. Obtained by the interaction of 60% oleum with methanol at 150°C. It is a methylating agent in organic synthesis. Carcinogen, affects the eyes, skin, respiratory organs.

^ Sodium thiosulfate Na 2 S 2 O 3

Salt of thiosulfuric acid, in which two sulfur atoms have different oxidation states: +6 and -2. Crystalline substance, highly soluble in water. It is produced in the form of Na 2 S 2 O 3 5H 2 O crystalline hydrate, commonly called hyposulfite. Obtained by the interaction of sodium sulfite with sulfur during boiling:

Na 2 SO 3 + S \u003d 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 \u003d 2H 2 SO 4 + 2NaCl + 6HCl

The use of sodium thiosulfate to absorb chlorine (in the first gas masks) was based on this reaction.

Sodium thiosulfate is oxidized somewhat differently by weak oxidizing agents. In this case, salts of tetrathionic acid are formed, for example:

2Na 2 S 2 O 3 + I 2 \u003d Na 2 S 4 O 6 + 2NaI

Sodium thiosulfate is a by-product in the production of NaHSO 3 , sulfur dyes, in the purification of industrial gases from sulfur. It is used to remove traces of chlorine after bleaching fabrics, to extract silver from ores; is a fixer in photography, a reagent in iodometry, an antidote for poisoning with arsenic, mercury compounds, an anti-inflammatory agent.

Sulfur(IV) oxide and sulfurous acid

Sulfur oxide (IV), or sulfur dioxide, under normal conditions, a colorless gas with a pungent suffocating odor. When cooled to -10°C, it liquefies into a colorless liquid.

Receipt

1. Under laboratory conditions, sulfur oxide (IV) is obtained from salts of sulfurous acid by the action of strong acids on them:

Na 2 SO 3 + H 2 SO 4 \u003d Na 2 SO 4 + S0 2 + H 2 O 2NaHSO 3 + H 2 SO 4 \u003d Na 2 SO 4 + 2SO 2 + 2H 2 O 2HSO - 3 + 2H + \u003d 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 \u003d CuSO 4 + SO 2 + 2H 2 O

Cu + 4Н + + 2SO 2- 4 \u003d Cu 2+ + SO 2- 4 + SO 2 + 2H 2 O

3. Sulfur oxide (IV) is also formed when sulfur is burned in air or oxygen:

4. Under industrial conditions, SO 2 is obtained by roasting pyrite FeS 2 or sulfurous ores of non-ferrous metals (zinc blende ZnS, lead luster PbS, etc.):

4FeS 2 + 11O 2 \u003d 2Fe 2 O 3 + 8SO 2

Structural formula of the SO 2 molecule:

Four sulfur electrons and four electrons from two oxygen atoms take part in the formation of bonds in the SO 2 molecule. The mutual repulsion of the bonding electron pairs and the unshared electron pair of sulfur gives the molecule an angular shape.

Chemical properties

1. Sulfur oxide (IV) exhibits all the properties of acidic oxides:

Interaction with water

Interaction with alkalis,

Interaction with basic oxides.

2. Sulfur oxide (IV) is characterized by reducing properties:

S +4 O 2 + O 0 2 "2S +6 O -2 3 (in the presence of a catalyst, when heated)

But in the presence of strong reducing agents, SO 2 behaves like an oxidizing agent:

Sulfur oxide (IV) is highly soluble in water (40 volumes of SO 2 are dissolved in 1 volume at 20 ° C). In this case, sulfurous acid exists only in an aqueous solution:

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 removing SO 2 (blowing through a nitrogen solution or heating), the reaction proceeds towards the starting materials. Sulfuric acid solution always contains sulfur oxide (IV), which gives it a pungent odor.

Sulfurous acid has all the properties of acids. Dissociates in solution 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 \u003d Na 2 SO 3 + 2H 2 O

Hydrosulfites are obtained with a lack of alkali:

H 2 SO 3 + NaOH \u003d 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. So, under the action of oxygen, sulfites are oxidized to sulfates:

2Na 2 S +4 O 3 + O 0 2 \u003d 2Na 2 S +6 O -2 4

The oxidation of sulfurous acid with bromine and potassium permanganate proceeds even more easily:

5H 2 S +4 O 3 +2KMn +7 O 4 \u003d 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:

Salts of sulfurous acid dissolve almost all hydrosulfites and sulfites of alkali metals.

3. Since H 2 SO 3 is a weak acid, the action of acids on sulfites and hydrosulfites releases SO 2. This method is usually used when obtaining SO 2 in the laboratory:

NaHSO 3 + H 2 SO 4 \u003d 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 increases in the solution:

Na 2 SO 3 + NON "NaHSO 3 + NaOH

Application

Sulfur oxide (IV) and sulfurous acid decolorize many dyes, forming colorless compounds with them. The latter can decompose again when heated or in the light, as a result of which the color is restored. Therefore, the bleaching effect of SO 2 and H 2 SO 3 is different from the bleaching effect of chlorine. Usually, sulfur (IV) rxide whitens wool, silk and straw.

An important application is the solution of calcium hydrosulfite CaHSO 3 (sulfite liquor), which is used to treat wood and paper pulp.