Hydrogen sulfide - H2S
Sulfur compounds -2, +4, +6. Qualitative reactions to sulfides, sulfites, sulfates.
Interaction Receive:
1. hydrogen with sulfur at t - 300 0
2. when acting on sulfides of mineral acids:
Na 2 S + 2HCl \u003d 2 NaCl + H 2 S
Physical properties:
colorless gas, with the smell of rotten eggs, poisonous, heavier than air, dissolving in water, forms a weak hydrogen sulfide acid.
Chemical properties
Acid-base properties
1. A solution of hydrogen sulfide in water - hydrosulfide acid - is a weak dibasic acid, therefore it dissociates in steps:
H 2 S ↔ HS - + H +
HS - ↔ H - + S 2-
2. Hydrosulfuric acid has the general properties of acids, reacts with metals, basic oxides, bases, salts:
H 2 S + Ca \u003d CaS + H 2
H 2 S + CaO \u003d CaS + H 2 O
H 2 S + 2NaOH \u003d Na 2 S + 2H 2 O
H 2 S + CuSO 4 \u003d CuS ↓ + H 2 SO 4
All acidic salts - hydrosulfides - are highly soluble in water. Normal salts - sulfides - dissolve in water in different ways: sulfides of alkali and alkaline earth metals are highly soluble, sulfides of other metals are insoluble in water, and sulfides of copper, lead, mercury and some other heavy metals do not dissolve even in acids (except nitric acid)
CuS + 4HNO 3 \u003d Cu (NO 3) 2 + 3S + 2NO + 2H 2 O
Soluble sulfides undergo hydrolysis - at the anion.
Na 2 S ↔ 2Na + + S 2-
S 2- +HOH ↔HS - +OH -
Na 2 S + H 2 O ↔ NaHS + NaOH
A qualitative reaction to hydrosulfide acid and its soluble salts (i.e., to the sulfide ion S 2-) is their interaction with soluble lead salts, with the formation of a black PbS precipitate
Na 2 S + Pb (NO 3) 2 \u003d 2NaNO 3 + PbS ↓
Pb 2+ + S 2- = PbS↓
Shows only restorative properties, tk. the sulfur atom has the lowest oxidation state -2
1. with oxygen
a) lacking
2H 2 S -2 + O 2 0 \u003d S 0 + 2H 2 O -2
b) with excess oxygen
2H 2 S + 3O 2 \u003d 2SO 2 + 2H 2 O
2. with halogens (discoloration of bromine water)
H 2 S -2 + Br 2 \u003d S 0 + 2HBr -1
3. with conc. HNO3
H 2 S + 2HNO 3 (k) \u003d S + 2NO 2 + 2H 2 O
b) with strong oxidizing agents (KMnO 4, K 2 CrO 4 in an acidic environment)
2KMnO 4 + 3H 2 SO 4 + 5H 2 S \u003d 5S + 2MnSO 4 + K 2 SO 4 + 8H 2 O
c) hydrosulfide acid is oxidized not only by strong oxidizing agents, but also by weaker ones, for example, iron (III) salts, sulfurous acid, etc.
2FeCl 3 + H 2 S = 2FeCl 2 + S + 2HCl
H 2 SO 3 + 2H 2 S \u003d 3S + 3H 2 O
Receipt
1. combustion of sulfur in oxygen.
2. combustion of hydrogen sulfide in excess O 2
2H 2 S + 3O 2 \u003d 2SO 2 + 2H 2 O
3. sulfide oxidation
2CuS + 3O 2 \u003d 2SO 2 + 2CuO
4. interaction of sulfites with acids
Na 2 SO 3 + H 2 SO 4 \u003d Na 2 SO 4 + SO 2 + H 2 O
5. interaction of metals in a series of activities after (H 2) with conc. H2SO4
Cu + 2H 2 SO 4 \u003d CuSO 4 + SO 2 + 2H 2 O
Physical Properties
Gas, colorless, with a suffocating smell of burnt sulfur, poisonous, more than 2 times heavier than air, highly soluble in water (at room temperature, about 40 volumes of gas dissolve in one volume).
Chemical properties:
Acid-base properties
SO 2 is a typical acidic oxide.
1.with alkalis, forming two types of salts: sulfites and hydrosulfites
2KOH + SO 2 \u003d K 2 SO 3 + H 2 O
KOH + SO 2 \u003d KHSO 3 + H 2 O
2.with basic oxides
K 2 O + SO 2 \u003d K 2 SO 3
3. weak sulfurous acid is formed with water
H 2 O + SO 2 \u003d H 2 SO 3
Sulfurous acid exists only in solution, is a weak acid,
has all the common properties of acids.
4. qualitative reaction to sulfite - ion - SO 3 2 - action of mineral acids
Na 2 SO 3 + 2HCl \u003d 2Na 2 Cl + SO 2 + H 2 O smell of burnt sulfur
redox properties
In OVR, it can be both an oxidizing agent and a reducing agent, because the sulfur atom in SO 2 has an intermediate oxidation state of +4.
As an oxidizing agent:
SO 2 + 2H 2 S = 3S + 2H 2 S
As a restorer:
2SO 2 +O 2 \u003d 2SO 3
Cl 2 + SO 2 + 2H 2 O \u003d H 2 SO 4 + 2HCl
2KMnO 4 + 5SO 2 + 2H 2 O \u003d K 2 SO 4 + 2H 2 SO 4 + 2MnSO 4
Sulfur oxide (VI) SO 3 (sulfuric anhydride)
Receipt:
Sulfur dioxide oxidation
2SO 2 + O 2 = 2SO 3 ( t 0 , cat)
Physical Properties
A colorless liquid, at temperatures below 17 0 С it turns into a white crystalline mass. Thermally unstable compound, completely decomposes at 700 0 C. It is highly soluble in water, in anhydrous sulfuric acid and reacts with it to form oleum
SO 3 + H 2 SO 4 \u003d H 2 S 2 O 7
Chemical properties
Acid-base properties
A typical acidic oxide.
1.with alkalis, forming two types of salts: sulfates and hydrosulfates
2KOH + SO 3 \u003d K 2 SO 4 + H 2 O
KOH + SO 3 \u003d KHSO 4 + H 2 O
2.with basic oxides
CaO + SO 2 \u003d CaSO 4
3. with water
H 2 O + SO 3 \u003d H 2 SO 4
redox properties
Sulfur oxide (VI) - a strong oxidizing agent, usually reduced to SO 2
3SO 3 + H 2 S \u003d 4SO 2 + H 2 O
Sulfuric acid H 2 SO 4
Getting sulfuric acid
In industry, acid is produced by the contact method:
1. pyrite firing
4FeS 2 + 11O 2 \u003d 2Fe 2 O 3 + 8SO 2
2. oxidation of SO 2 to SO 3
2SO 2 + O 2 = 2SO 3 ( t 0 , cat)
3. dissolution of SO 3 in sulfuric acid
n SO 3 + H 2 SO 4 \u003d H 2 SO 4 ∙ n SO 3 (oleum)
H 2 SO 4 ∙ n SO 3 + H 2 O \u003d H 2 SO 4
Physical Properties
H 2 SO 4 is a heavy oily liquid, odorless and colorless, hygroscopic. Miscible with water in any ratio, when concentrated sulfuric acid is dissolved in water, a large amount of heat is released, so it must be carefully poured into water, and not vice versa (first water, then acid, otherwise big trouble will happen)
A solution of sulfuric acid in water with an H 2 SO 4 content of less than 70% is usually called dilute sulfuric acid, more than 70% is concentrated.
Chemical properties
Acid-base
Dilute sulfuric acid exhibits all the characteristic properties of strong acids. Dissociates in aqueous solution:
H 2 SO 4 ↔ 2H + + SO 4 2-
1. with basic oxides
MgO + H 2 SO 4 \u003d MgSO 4 + H 2 O
2. with bases
2NaOH + H 2 SO 4 \u003d Na 2 SO 4 + 2H 2 O
3. with salts
BaCl 2 + H 2 SO 4 \u003d BaSO 4 ↓ + 2HCl
Ba 2+ + SO 4 2- = BaSO 4 ↓ (white precipitate)
Qualitative reaction to the sulfate ion SO 4 2-
Due to the higher boiling point, compared to other acids, sulfuric acid displaces them from salts when heated:
NaCl + H 2 SO 4 \u003d HCl + NaHSO 4
redox properties
In dilute H 2 SO 4, oxidizing agents are H + ions, and in concentrated H 2 SO 4 - sulfate ions SO 4 2
In dilute sulfuric acid, metals that are in the order of activity up to hydrogen dissolve, while sulfates are formed and hydrogen is released
Zn + H 2 SO 4 \u003d ZnSO 4 + H 2
Concentrated sulfuric acid is a vigorous oxidizing agent, especially when heated. It oxidizes many metals, non-metals, inorganic and organic substances.
H 2 SO 4 (to) oxidizing agent S +6
With more active metals, sulfuric acid, depending on the concentration, can be reduced to a variety of products.
Zn + 2H 2 SO 4 \u003d ZnSO 4 + SO 2 + 2H 2 O
3Zn + 4H 2 SO 4 = 3ZnSO 4 + S + 4H 2 O
4Zn + 5H 2 SO 4 = 4ZnSO 4 + H 2 S + 4H 2 O
Concentrated sulfuric acid oxidizes some non-metals (sulfur, carbon, phosphorus, etc.), reducing to sulfur oxide (IV)
S + 2H 2 SO 4 \u003d 3SO 2 + 2H 2 O
C + 2H 2 SO 4 \u003d 2SO 2 + CO 2 + 2H 2 O
Interaction with some complex substances
H 2 SO 4 + 8HI \u003d 4I 2 + H 2 S + 4 H 2 O
H 2 SO 4 + 2HBr \u003d Br 2 + SO 2 + 2H 2 O
Salts of sulfuric acid
2 types of salts: sulfates and hydrosulfates
Salts of sulfuric acid have all the common properties of salts. Their relation to heating is special. Sulfates of active metals (Na, K, Ba) do not decompose even when heated above 1000 0 C, salts of less active metals (Al, Fe, Cu) decompose even with slight heating
In 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 does the oxidizing agent SO 2 react with stronger reducing agents, for example with:
SO 2 + 2H 2 S \u003d 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.:
2SO 2 + O 2 \u003d 2SO 3
SO 2 + Cl 2 + 2H 2 O \u003d H 2 SO 3 + 2HCl
Receipt
1) Sulfur dioxide is formed during the combustion of sulfur:
2) In industry, it is obtained by firing pyrite:
3) In the laboratory, sulfur dioxide can be obtained:
Cu + 2H 2 SO 4 \u003d 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. 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 ° C turns into a white crystalline mass. It absorbs moisture very well (hygroscopic).
Chemical properties
Acid-base properties
How a typical acid oxide sulfuric anhydride interacts:
SO 3 + CaO = CaSO 4
c) with water:
SO 3 + H 2 O \u003d 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.
Oleum formation: H 2 SO 4 + n SO 3 \u003d 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 \u003d 4SO 2 + H 2 O
Getting and using
Sulfuric anhydride is formed during the oxidation of sulfur dioxide:
2SO 2 + O 2 \u003d 2SO 3
In its pure form, sulfuric anhydride has no practical value. It is obtained as an intermediate 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 \u003d Fe 2 O 3 + SO 3 + SO 2 or a mixture with: 6KNO 3 + 5S \u003d 3K 2 SO 4 + 2SO 3 + 3N 2, and the emitted vapors of sulfuric anhydride were condensed. Absorbing moisture, they turned into oleum. Depending on the method of preparation, H 2 SO 4 was called vitriol oil or sulfur oil. In 1595, the alchemist Andreas Libavius established the identity of both substances.
For a long time, vitriol oil was not widely used. Interest in it greatly increased after the 18th century. Indigo carmine, a stable blue dye, 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 with sulfur was burned in the upper part of the chamber, then air was let in there. 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, nitric acid was used (it gives when decomposed in the chamber). To return nitrous gases to the system, special towers were designed, which gave the name to the whole process - the tower process. Factories operating according to 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 water (and not vice versa!) And mix the solution.
A solution of sulfuric acid in water with an H2SO4 content of less than 70% is usually called dilute sulfuric acid, and a solution of more than 70% is called 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 \u003d Na 2 SO 4 + 2H 2 O
H 2 SO 4 + BaCl 2 \u003d BaSO 4 ↓ + 2HCl
The process of interaction of Ba 2+ ions with sulfate ions SO 4 2+ leads to the formation of a white insoluble precipitate BaSO 4 . it qualitative reaction to sulfate ion.
Redox properties
In dilute H 2 SO 4 , H + ions are oxidizing agents, and in concentrated H 2 SO 4 sulfate ions are SO 4 2+ . SO 4 2+ ions are stronger oxidizing agents than H + ions (see diagram). 
AT dilute sulfuric acid dissolve metals that are in the electrochemical series of voltages to hydrogen. In this case, metal sulfates are formed and released:
Zn + H 2 SO 4 \u003d ZnSO 4 + H 2
Metals that are in the electrochemical series of voltages after hydrogen 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 in the electrochemical series of voltages after hydrogen (Cu, Ag, Hg), metal sulfates are formed, as well as the product of sulfuric acid reduction - SO 2.
Reaction of sulfuric acid with zinc With more active metals (Zn, Al, Mg), concentrated sulfuric acid can be reduced to free. For example, when sulfuric acid interacts with, depending on the acid concentration, various products of sulfuric acid reduction can be formed simultaneously - SO 2, S, H 2 S:
Zn + 2H 2 SO 4 \u003d 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 therefore it is transported in iron tanks:
Fe + H 2 SO 4 ≠
Concentrated sulfuric acid oxidizes some non-metals (, etc.), recovering to sulfur oxide (IV) SO 2:
S + 2H 2 SO 4 \u003d 3SO 2 + 2H 2 O
C + 2H 2 SO 4 \u003d 2SO 2 + CO 2 + 2H 2 O
Getting and using
In industry, sulfuric acid is obtained by contact. The acquisition process takes place in three stages:
- Obtaining SO 2 by roasting pyrite:
4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8SO 2
- Oxidation of SO 2 to SO 3 in the presence of a catalyst - vanadium (V) oxide:
2SO 2 + O 2 \u003d 2SO 3
- Dissolution of SO 3 in sulfuric acid:
H2SO4+ n SO 3 \u003d 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 pouring it into water. This can be expressed in a diagram:
H 2 SO 4 ∙ n SO 3 + H 2 O \u003d H 2 SO 4
Sulfuric acid finds various applications in various fields 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.
Salts of sulfuric acid
Most sulfates are highly soluble in water (slightly soluble CaSO 4 , even less PbSO 4 and practically insoluble BaSO 4). Some sulfates containing water of crystallization are called vitriol:
CuSO 4 ∙ 5H 2 O copper sulfate
FeSO 4 ∙ 7H 2 O ferrous sulfate
Salts of sulfuric acid have everything. Their relation to heating is special.
Sulfates of active metals ( , ) do not decompose even at 1000 ° C, while others (Cu, Al, Fe) - decompose upon slight heating into metal oxide and SO 3:
CuSO 4 \u003d CuO + SO 3
Download:
Download free abstract on the topic: "Production of sulfuric acid by contact method"
You can download essays on other topics
*on the image of the record is a photograph of copper sulphate
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 of 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 obtaining sulfur from natural compounds, for example, from sulfur pyrite, 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 hard brittle lemon-yellow substance. 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 It is 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 incomplete external energy level, can attach two electrons and show 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 mainly 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 two p-electrons of the outer level of the sulfur atom are involved in the formation of H-S-H bonds. 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 unpleasant smell of sewage, 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. It must be handled in fume hoods or hermetically sealed appliances. The 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 of nitric acid, free sulfur is formed.
6. A solution of hydrogen sulfide has an acidic reaction due to dissociations:
H 2 SH + +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, alkali metal sulfides, when dissolved in water, have an alkaline reaction:
Na 2 S+HOHNaHS+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 OH 2 SO 3
The reaction is reversible. In an aqueous solution, sulfur oxide (IV) and sulfurous acid are in chemical equilibrium, which can be shifted. When binding H 2 SO 3 (neutralization of acid
You) the reaction proceeds towards the formation of sulfurous acid; when 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 + NOHNaHSO 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.
An important application is the 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. The most important is the contact method of obtaining
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. It absorbs moisture with the release of a large amount of heat, so you can not add water to concentrated sulfuric acid - the 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 of 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 amount of it is spent 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 clean oil 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 is a blue toxic substance. 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 fixative in photography, a reagent in iodometry, an antidote for poisoning with arsenic, mercury compounds, and an anti-inflammatory agent.
Part I
1. Hydrogen sulfide.
1) The structure of the molecule:
2) Physical properties: colorless gas, with a pungent smell of rotten eggs, heavier than air.
3) Chemical properties (finish the reaction equations and consider the equations in the light of TED or from the standpoint of redox).
4) Hydrogen sulfide in nature: in the form of compounds - sulfides, in a free form - in volcanic gases.
2. Sulfur oxide (IV) - SO2
1) Getting in the industry. Write down the reaction equations and consider them in terms of oxidation-reduction.
2) Obtaining in the laboratory. Write down the reaction equation and consider it in the light of TED:
3) Physical properties: gas with a pungent, suffocating odour.
4) Chemical properties.
3. Sulfur oxide (VI) - SO3.
1) Obtaining by synthesis from sulfur oxide (IV):
2) Physical properties: liquid, heavier than water, mixed with sulfuric acid - oleum.
3) Chemical properties. Shows typical properties of acidic oxides:
Part II
1. Describe 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) burning
2. Describe the reaction of the interaction of sulfur oxide (IV) 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 oxide (IV) 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 from the standpoint of oxidation-reduction.
6. Write down the equations for the reactions of transitions, deciphering the unknown formulas:
7. Write a cinquain 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 report on the toxicity of hydrogen sulfide (pay attention to its characteristic smell!) And first aid in case of poisoning with this gas. Write down the message plan in a special notebook.
hydrogen sulfide
A colorless gas with a rotten egg odor. It is found in the air by smell, even in small concentrations. In nature, it is found in the water of mineral springs, seas, volcanic gases. It is formed during the decomposition of proteins in the absence of oxygen. It can be released into the air in a number of chemical and textile industries, during the extraction and processing of oil, from sewage.
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. Upon termination of exposure, 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 danger of poisoning increases due to loss of smell. Cardiac weakness and respiratory failure, coma gradually increase.
First aid - removal of the victim from the polluted atmosphere, inhalation of oxygen, artificial respiration; means that excite the respiratory center, warming the body. Glucose, vitamins, iron preparations are also recommended.
Prevention - sufficient ventilation, sealing of some production operations. When descending workers into wells and containers containing hydrogen sulfide, they must use gas masks and life belts on ropes. Gas rescue service is obligatory in mines, in places of extraction and at oil refineries.
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:
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 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.
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.
An important application is the solution of calcium hydrosulfite CaHSO 3 (sulfite liquor), which is used to treat wood and paper pulp.
