Sulfur oxide (sulfur dioxide, sulfur dioxide, sulfur dioxide) is a colorless gas that under normal conditions has a sharp characteristic odor (similar to the smell of a burning match). It liquefies under pressure at room temperature. Sulfur dioxide is soluble in water, and unstable sulfuric acid is formed. This substance is also soluble in sulfuric acid and ethanol. This is one of the main components that make up volcanic gases.

1. Sulfur dioxide dissolves in water, resulting in sulfurous acid. Under normal conditions, this reaction is reversible.

SO2 (sulfur dioxide) + H2O (water) = H2SO3 (sulfurous acid).

2. With alkalis, sulfur dioxide forms sulfites. For example: 2NaOH (sodium hydroxide) + SO2 (sulfur dioxide) = Na2SO3 (sodium sulfite) + H2O (water).

3. The chemical activity of sulfur dioxide is quite high. The reducing properties of sulfur dioxide are most pronounced. In such reactions, the oxidation state of sulfur increases. For example: 1) SO2 (sulfur dioxide) + Br2 (bromine) + 2H2O (water) = H2SO4 (sulfuric acid) + 2HBr (hydrogen bromide); 2) 2SO2 (sulfur dioxide) + O2 (oxygen) = 2SO3 (sulfite); 3) 5SO2 (sulfur dioxide) + 2KMnO4 (potassium permanganate) + 2H2O (water) = 2H2SO4 (sulfuric acid) + 2MnSO4 (manganese sulfate) + K2SO4 (potassium sulfate).

The last reaction is an example of a qualitative reaction to SO2 and SO3. The solution becomes purple in color.)

4. In the presence of strong reducing agents, sulfur dioxide can exhibit oxidizing properties. For example, in order to extract sulfur from exhaust gases in the metallurgical industry, they use the reduction of sulfur dioxide with carbon monoxide (CO): SO2 (sulfur dioxide) + 2CO (carbon monoxide) = 2CO2 + S (sulfur).

Also, the oxidizing properties of this substance are used to obtain phosphorous acid: PH3 (phosphine) + SO2 (sulfur dioxide) = H3PO2 (phosphoric acid) + S (sulfur).

Where is sulfur dioxide used?

Sulfur dioxide is mainly used to produce sulfuric acid. It is also used in the production of low-alcohol drinks (wine and other mid-price drinks). Due to the property of this gas to kill various microorganisms, it is used to fumigate warehouses and vegetable stores. In addition, sulfur oxide is used to bleach wool, silk, and straw (those materials that cannot be bleached with chlorine). In laboratories, sulfur dioxide is used as a solvent and in order to obtain various salts of sulfur dioxide.

Physiological effects

Sulfur dioxide has strong toxic properties. Symptoms of poisoning are cough, runny nose, hoarseness, a peculiar taste in the mouth, and severe sore throat. When sulfur dioxide is inhaled in high concentrations, difficulty swallowing and choking, speech disturbance, nausea and vomiting occur, and acute pulmonary edema may develop.

MPC of sulfur dioxide:
- indoors - 10 mg/m³;
- average daily maximum one-time exposure in atmospheric air - 0.05 mg/m³.

Sensitivity to sulfur dioxide varies among individuals, plants, and animals. For example, among trees the most resistant are oak and birch, and the least resistant are spruce and pine.

DEFINITION

Hydrogen sulfide is a colorless gas with a characteristic odor of rotting protein.

It is slightly heavier than air, liquefies at a temperature of -60.3 o C and solidifies at -85.6 o C. In air, hydrogen sulfide burns with a bluish flame, forming sulfur dioxide and water:

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

If you introduce some cold object, such as a porcelain cup, into the hydrogen sulfide flame, the temperature of the flame drops significantly and the hydrogen sulfide oxidizes only to free sulfur, which settles on the cup in the form of a yellow coating:

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

Hydrogen sulfide is highly flammable; its mixture with air explodes. Hydrogen sulfide is very poisonous. Prolonged inhalation of air containing this gas, even in small quantities, causes severe poisoning.

At 20 o C, one volume of water dissolves 2.5 volumes of hydrogen sulfide. A solution of hydrogen sulfide in water is called hydrogen sulfide water. When standing in the air, especially in the light, hydrogen sulfide water soon becomes cloudy from the sulfur released. This occurs as a result of the oxidation of hydrogen sulfide by atmospheric oxygen.

Production of hydrogen sulfide

At high temperatures, sulfur reacts with hydrogen to form hydrogen sulfide gas.

In practice, hydrogen sulfide is usually produced by the action of dilute acids on sulfur metals, for example iron sulfide:

FeS + 2HCl = FeCl 2 + H 2 S.

More pure hydrogen sulfide can be obtained by hydrolysis of CaS, BaS or A1 2 S 3. The purest gas is obtained by the direct reaction of hydrogen and sulfur at 600 °C.

Chemical properties of hydrogen sulfide

A solution of hydrogen sulfide in water has the properties of an acid. Hydrogen sulfide is a weak dibasic acid. It dissociates step by step and mainly according to the first step:

H 2 S↔H + + HS - (K 1 = 6 × 10 -8).

Second stage dissociation

HS - ↔H + + S 2- (K 2 = 10 -14)

occurs to a negligible extent.

Hydrogen sulfide is a strong reducing agent. When exposed to strong oxidizing agents, it is oxidized to sulfur dioxide or sulfuric acid; the depth of oxidation depends on the conditions: temperature, pH of the solution, concentration of the oxidizing agent. For example, the reaction with chlorine usually proceeds to form sulfuric acid:

H 2 S + 4Cl 2 + 4H 2 O = H 2 SO 4 + 8HCl.

Medium salts of hydrogen sulfide are called sulfides.

Application of hydrogen sulfide

The use of hydrogen sulfide is quite limited, which is primarily due to its high toxicity. It has found application in laboratory practice as a precipitant for heavy metals. Hydrogen sulfide serves as a raw material for the production of sulfuric acid, sulfur in elemental form and sulfides

Examples of problem solving

EXAMPLE 1

Exercise	Determine how many times heavier than air is hydrogen sulfide H 2 S.
Solution	The ratio of the mass of a given gas to the mass of another gas taken in the same volume, at the same temperature and the same pressure is called the relative density of the first gas to the second. This value shows how many times the first gas is heavier or lighter than the second gas. The relative molecular weight of air is taken to be 29 (taking into account the content of nitrogen, oxygen and other gases in the air). It should be noted that the concept of “relative molecular mass of air” is used conditionally, since air is a mixture of gases. D air (H 2 S) = M r (H 2 S) / M r (air); D air (H 2 S) = 34 / 29 = 1.17. M r (H 2 S) = 2 × A r (H) + A r (S) = 2 × 1 + 32 = 2 + 32 = 34.
Answer	Hydrogen sulfide H 2 S is 1.17 times heavier than air.

EXAMPLE 2

Exercise	Find the hydrogen density of a mixture of gases in which the volume fraction of oxygen is 20%, hydrogen is 40%, and the rest is hydrogen sulfide H 2 S.
Solution	The volume fractions of gases will coincide with the molar ones, i.e. with fractions of quantities of substances, this is a consequence of Avogadro’s law. Let's find the conditional molecular weight of the mixture: M r conditional (mixture) = φ (O 2) × M r (O 2) + φ (H 2) × M r (H 2) + φ (H 2 S) × M r (H 2 S);

- (hydrogen sulfide) H2S, a colorless gas with the smell of rotten eggs; melting point?85.54.C, boiling point?60.35.C; at 0.C it liquefies under a pressure of 1 MPa. Reducing agent. A by-product during the refining of petroleum products, coking of coal, etc.; formed during decomposition... ... Big Encyclopedic Dictionary

HYDROGEN Sulfide- (H2S), a colorless, poisonous gas with the smell of rotten eggs. Formed during decay processes, found in crude oil. Obtained by the action of sulfuric acid on metal sulfides. Used in traditional QUALITATIVE ANALYSIS. Properties: temperature... ... Scientific and technical encyclopedic dictionary

HYDROGEN Sulfide- HYDROGEN Sulfide, hydrogen sulfide, many others. no, husband (chem.). A gas produced by the rotting of protein substances, giving off the smell of rotten eggs. Ushakov's explanatory dictionary. D.N. Ushakov. 1935 1940 … Ushakov's Explanatory Dictionary

HYDROGEN Sulfide- HYDROGEN SULFIDE, huh, husband. A colorless gas with a sharp, unpleasant odor, formed during the decomposition of protein substances. | adj. hydrogen sulfide, oh, oh. Ozhegov's explanatory dictionary. S.I. Ozhegov, N.Yu. Shvedova. 1949 1992 … Ozhegov's Explanatory Dictionary

hydrogen sulfide- noun, number of synonyms: 1 gas (55) ASIS Dictionary of Synonyms. V.N. Trishin. 2013… Synonym dictionary

HYDROGEN Sulfide- colorless poisonous gas H2S with an unpleasant specific odor. It has slightly acidic properties. 1 liter of C. at t 0 °C and a pressure of 760 mm is 1.539 g. It is found in oils, natural waters, and gases of biochemical origin, such as... ... Geological encyclopedia

HYDROGEN Sulfide- HYDROGEN Sulfide, H2S (molecular weight 34.07), a colorless gas with a characteristic odor of rotten eggs. A liter of gas under normal conditions (0°, 760 mm) weighs 1.5392 g. Boiling temperature 62°, melting 83°; S. is part of gaseous emissions... ... Great Medical Encyclopedia

hydrogen sulfide- - Topics of biotechnology EN hydrogen sulfide ... Technical Translator's Guide

hydrogen sulfide- HYDROGEN SULFIDE, a, m A colorless gas with a sharp, unpleasant odor, formed during the decomposition of protein substances and representing a compound of sulfur with hydrogen. Hydrogen sulfide is found in some mineral waters and medicinal muds and is used... ... Explanatory dictionary of Russian nouns

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O.S.ZAYTSEV

CHEMISTRY BOOK

FOR SECONDARY SCHOOL TEACHERS,
STUDENTS OF PEDAGOGICAL UNIVERSITIES AND SCHOOLCHILDREN OF 9–10 GRADES,
WHO DECIDED TO DEVOTE THEMSELVES TO CHEMISTRY AND NATURAL SCIENCE

TEXTBOOK TASK LABORATORY PRACTICAL SCIENTIFIC STORIES FOR READING

Continuation. See No. 4–14, 16–28, 30–34, 37–44, 47, 48/2002;
1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 12, 13, 14, 15, 16, 17, 18, 19, 20, 21, 22, 23,
24, 25-26, 27-28, 29, 30, 31, 32, 35, 36, 37, 39, 41, 42, 43, 44 , 46, 47/2003;
1, 2, 3, 4, 5, 7, 11, 13, 14, 16, 17, 20, 22, 24/2004

§ 8.1. Redox reactions

LABORATORY RESEARCH
(continuation)

2. Ozone is an oxidizing agent.

Ozone is the most important substance for nature and humans.

Ozone creates an ozonosphere around the Earth at an altitude of 10 to 50 km with a maximum ozone content at an altitude of 20–25 km. Being in the upper layers of the atmosphere, ozone does not allow most of the sun's ultraviolet rays, which have a detrimental effect on humans, animals and plants, to reach the Earth's surface. In recent years, areas of the ozonosphere with greatly reduced ozone content, the so-called ozone holes, have been discovered. It is not known whether ozone holes have formed before. The reasons for their occurrence are also unclear. It is assumed that chlorine-containing freons from refrigerators and perfume cans, under the influence of ultraviolet radiation from the Sun, release chlorine atoms, which react with ozone and thereby reduce its concentration in the upper layers of the atmosphere. Scientists are extremely concerned about the danger of ozone holes in the atmosphere.
In the lower layers of the atmosphere, ozone is formed as a result of a series of sequential reactions between atmospheric oxygen and nitrogen oxides emitted by poorly adjusted car engines and discharges from high-voltage power lines. Ozone is very harmful to breathing - it destroys the tissue of the bronchi and lungs. Ozone is extremely toxic (more powerful than carbon monoxide). The maximum permissible concentration in the air is 10–5%.
Thus, ozone in the upper and lower layers of the atmosphere has opposite effects on humans and the animal world.
Ozone, along with chlorine, is used to treat water to break down organic impurities and kill bacteria. However, both chlorination and ozonation of water have their advantages and disadvantages. When water is chlorinated, bacteria are almost completely destroyed, but organic substances of a carcinogenic nature that are harmful to health (promote the development of cancer) are formed - dioxins and similar compounds. When water is ozonized, such substances are not formed, but ozone does not kill all bacteria, and after some time the remaining living bacteria multiply abundantly, absorbing the remains of killed bacteria, and the water becomes even more contaminated with bacterial flora. Therefore, ozonation of drinking water is best used when it is used quickly. Ozonation of water in swimming pools is very effective when water continuously circulates through the ozonizer. Ozone is also used for air purification. It is one of the environmentally friendly oxidizing agents that do not leave harmful products of its decomposition.
Ozone oxidizes almost all metals except gold and platinum group metals.

Chemical methods for producing ozone are ineffective or too dangerous. Therefore, we advise you to obtain ozone mixed with air in an ozonizer (the effect of a weak electrical discharge on oxygen) available in the school physics laboratory.

Ozone is most often obtained by acting on gaseous oxygen with a quiet electrical discharge (without glow or sparks), which occurs between the walls of the internal and external vessels of the ozonizer. The simplest ozonizer can be easily made from glass tubes with stoppers. You will understand how to do this from Fig. 8.4. The inner electrode is a metal rod (long nail), the outer electrode is a wire spiral. Air can be blown out with an aquarium air pump or a rubber bulb from a spray bottle. In Fig. 8.4 The inner electrode is located in a glass tube ( Why do you think?), but you can assemble an ozonizer without it. Rubber plugs are quickly corroded by ozone.

It is convenient to obtain high voltage from the induction coil of the car's ignition system by continuously opening the connection to a low voltage source (battery or 12 V rectifier).
The ozone yield is several percent.

Ozone can be detected qualitatively using a starch solution of potassium iodide. A strip of filter paper can be soaked in this solution, or the solution can be added to ozonized water, and air with ozone can be passed through the solution in a test tube. Oxygen does not react with iodide ion.
Reaction equation:

2I – + O 3 + H 2 O = I 2 + O 2 + 2OH – .

Write the equations for the reactions of electron gain and loss.
Bring a strip of filter paper moistened with this solution to the ozonizer. (Why should a potassium iodide solution contain starch?) Hydrogen peroxide interferes with the determination of ozone using this method. (Why?).
Calculate the EMF of the reaction using the electrode potentials:

3. Reducing properties of hydrogen sulfide and sulfide ion.

Hydrogen sulfide is a colorless gas with the smell of rotten eggs (some proteins contain sulfur).
To conduct experiments with hydrogen sulfide, you can use gaseous hydrogen sulfide, passing it through a solution with the substance being studied, or add pre-prepared hydrogen sulfide water to the solutions under study (this is more convenient). Many reactions can be carried out with a solution of sodium sulfide (reactions with the sulfide ion S 2–).
Work with hydrogen sulfide only under draft! Mixtures of hydrogen sulfide with air burn explosively.

Hydrogen sulfide is usually produced in a Kipp apparatus by reacting 25% sulfuric acid (diluted 1:4) or 20% hydrochloric acid (diluted 1:1) on iron sulfide in the form of pieces 1–2 cm in size. Reaction equation:

FeS (cr.) + 2H + = Fe 2+ + H 2 S (g.).

Small quantities of hydrogen sulfide can be obtained by placing crystalline sodium sulfide in a stoppered flask through which a dropping funnel with a stopcock and an outlet tube are passed. Slowly pouring 5–10% hydrochloric acid from the funnel (why not sulfur?), the flask is constantly shaken by shaking to avoid local accumulation of unreacted acid. If this is not done, unexpected mixing of components can lead to a violent reaction, expulsion of the stopper and destruction of the flask.
A uniform flow of hydrogen sulfide is obtained by heating hydrogen-rich organic compounds, such as paraffin, with sulfur (1 part paraffin to 1 part sulfur, 300 ° C).
To obtain hydrogen sulfide water, hydrogen sulfide is passed through distilled (or boiled) water. About three volumes of hydrogen sulfide gas dissolve in one volume of water. When standing in air, hydrogen sulfide water gradually becomes cloudy. (Why?).
Hydrogen sulfide is a strong reducing agent: it reduces halogens to hydrogen halides, and sulfuric acid to sulfur dioxide and sulfur.
Hydrogen sulfide is poisonous. The maximum permissible concentration in the air is 0.01 mg/l. Even at low concentrations, hydrogen sulfide irritates the eyes and respiratory tract and causes headaches. Concentrations above 0.5 mg/l are life-threatening. At higher concentrations, the nervous system is affected. Inhaling hydrogen sulfide may cause cardiac and respiratory arrest. Sometimes hydrogen sulfide accumulates in caves and sewer wells, and a person trapped there instantly loses consciousness and dies.
At the same time, hydrogen sulfide baths have a healing effect on the human body.

3a. Reaction of hydrogen sulfide with hydrogen peroxide.

Study the effect of hydrogen peroxide solution on hydrogen sulfide water or sodium sulfide solution.
Based on the results of the experiments, compose reaction equations. Calculate the EMF of the reaction and draw a conclusion about the possibility of its passage.

3b. Reaction of hydrogen sulfide with sulfuric acid.

Pour concentrated sulfuric acid dropwise into a test tube with 2–3 ml of hydrogen sulfide water (or sodium sulfide solution). (carefully!) until turbidity appears. What is this substance? What other products might be produced in this reaction?
Write the reaction equations. Calculate the EMF of the reaction using the electrode potentials:

4. Sulfur dioxide and sulfite ion.

Sulfur dioxide, sulfur dioxide, is the most important atmospheric pollutant emitted by automobile engines when using poorly purified gasoline and by furnaces in which sulfur-containing coals, peat or fuel oil are burned. Every year, millions of tons of sulfur dioxide are released into the atmosphere due to the burning of coal and oil.
Sulfur dioxide occurs naturally in volcanic gases. Sulfur dioxide is oxidized by atmospheric oxygen into sulfur trioxide, which, absorbing water (vapor), turns into sulfuric acid. Falling acid rain destroys cement parts of buildings, architectural monuments, and sculptures carved from stone. Acid rain slows down the growth of plants and even leads to their death, and kills living organisms in water bodies. Such rains wash out phosphorus fertilizers, which are poorly soluble in water, from arable lands, which, when released into water bodies, lead to rapid proliferation of algae and rapid swamping of ponds and rivers.
Sulfur dioxide is a colorless gas with a pungent odor. Sulfur dioxide should be obtained and worked with under draft.

Sulfur dioxide can be obtained by placing 5–10 g of sodium sulfite in a flask closed with a stopper with an outlet tube and a dropping funnel. From a dropping funnel with 10 ml concentrated sulfuric acid (extreme caution!) pour it drop by drop onto the sodium sulfite crystals. Instead of crystalline sodium sulfite, you can use its saturated solution.
Sulfur dioxide can also be produced by the reaction between copper metal and sulfuric acid. In a round-bottomed flask equipped with a stopper with a gas outlet tube and a dropping funnel, place copper shavings or pieces of wire and pour a little sulfuric acid from the dropping funnel (about 6 ml of concentrated sulfuric acid is taken per 10 g of copper). To start the reaction, warm the flask slightly. After this, add the acid drop by drop. Write the equations for accepting and losing electrons and the total equation.
The properties of sulfur dioxide can be studied by passing the gas through a reagent solution, or in the form of an aqueous solution (sulfurous acid). The same results are obtained when using acidified solutions of sodium sulfites Na 2 SO 3 and potassium sulfites K 2 SO 3. Up to forty volumes of sulfur dioxide are dissolved in one volume of water (a ~6% solution is obtained).
Sulfur dioxide is toxic. With mild poisoning, a cough begins, a runny nose, tears appear, and dizziness begins. Increasing the dose leads to respiratory arrest.

4a. Interaction of sulfurous acid with hydrogen peroxide.

Predict the reaction products of sulfurous acid and hydrogen peroxide. Test your assumption with experience.
Add the same amount of 3% hydrogen peroxide solution to 2–3 ml of sulfurous acid. How to prove the formation of the expected reaction products?
Repeat the same experiment with acidified and alkaline solutions of sodium sulfite.
Write the reaction equations and calculate the emf of the process.
Select the electrode potentials you need:

4b. Reaction between sulfur dioxide and hydrogen sulfide.

This reaction takes place between gaseous SO 2 and H 2 S and serves to produce sulfur. The reaction is also interesting because the two air pollutants mutually destroy each other. Does this reaction take place between solutions of hydrogen sulfide and sulfur dioxide? Answer this question with experience.
Select electrode potentials to determine whether a reaction can occur in solution:

Try to carry out a thermodynamic calculation of the possibility of reactions. The thermodynamic characteristics of substances to determine the possibility of a reaction between gaseous substances are as follows:

In which state of substances - gaseous or in solution - are reactions more preferable?

Sulfur– element of the 3rd period and VIA group of the Periodic System, serial number 16, refers to chalcogens. The electronic formula of the atom is [ 10 Ne]3s 2 3p 4, the characteristic oxidation states are 0, -II, +IV and +VI, the S VI state is considered stable.

Scale of sulfur oxidation states:

The electronegativity of sulfur is 2.60 and is characterized by non-metallic properties. In hydrogen and oxygen compounds it is found in various anions and forms oxygen-containing acids and their salts, binary compounds.

In nature - fifteenth element by chemical abundance (seventh among non-metals). It is found in free (native) and bound form. A vital element for higher organisms.

Sulfur S. Simple substance. Yellow crystalline (α‑rhombic and β‑monoclinic,

at 95.5 °C) or amorphous (plastic). At the nodes of the crystal lattice there are S 8 molecules (non-planar rings of the “crown” type), amorphous sulfur consists of S n chains. A low-melting substance, the viscosity of the liquid passes through a maximum at 200 °C (breakdown of S 8 molecules, interweaving of S n chains). The pair contains molecules S 8, S 6, S 4, S 2. At 1500 °C, monoatomic sulfur appears (in chemical equations, for simplicity, any sulfur is depicted as S).

Sulfur is insoluble in water and under normal conditions does not react with it; it is highly soluble in carbon disulfide CS 2.

Sulfur, especially powdered sulfur, is highly active when heated. Reacts as an oxidizing agent with metals and non-metals:

but as reducing agent– with fluorine, oxygen and acids (boiling):

Sulfur undergoes dismutation in alkali solutions:

3S 0 + 6KOH (conc.) = 2K 2 S ‑II + K 2 S IV O 3 + 3H 2 O

At high temperatures (400 °C), sulfur displaces iodine from hydrogen iodide:

S + 2HI (g) = I 2 + H 2 S,

but in solution the reaction goes in the opposite direction:

I 2 + H 2 S (p) = 2 HI + S↓

Receipt: V industry smelted from natural deposits of native sulfur (using water vapor), released during desulfurization of coal gasification products.

Sulfur is used for the synthesis of carbon disulfide, sulfuric acid, sulfur (vat) dyes, in the vulcanization of rubber, as a means of protecting plants from powdery mildew, and for the treatment of skin diseases.

Hydrogen sulfide H 2 S. Anoxic acid. A colorless gas with a suffocating odor, heavier than air. The molecule has the structure of a doubly incomplete tetrahedron [::S(H) 2 ]

(sp 3 -hybridization, valet angle H – S–H is far from tetrahedral). Unstable when heated above 400 °C. Slightly soluble in water (2.6 l/1 l H 2 O at 20 °C), saturated decimolar solution (0.1 M, “hydrogen sulfide water”). A very weak acid in solution, practically does not dissociate in the second stage to S 2‑ ions (the maximum concentration of S 2‑ is 1 10 ‑ 13 mol/l). When exposed to air, the solution becomes cloudy (the inhibitor is sucrose). Neutralized by alkalis, but not completely by ammonia hydrate. Strong reducing agent. Enters into ion exchange reactions. A sulfiding agent precipitates differently colored sulfides with very low solubility from solution.

Qualitative reactions– precipitation of sulfides, as well as incomplete combustion of H 2 S with the formation of a yellow sulfur deposit on a cold object brought into the flame (porcelain spatula). A by-product of oil, natural and coke oven gas refining.

It is used in the production of sulfur, inorganic and organic sulfur-containing compounds as an analytical reagent. Extremely poisonous. Equations of the most important reactions:

Receipt: V industry– direct synthesis:

H 2 + S = H2S(150–200 °C)

or by heating sulfur with paraffin;

V laboratories– displacement from sulfides with strong acids

FeS + 2НCl (conc.) = FeCl 2 + H2S

or complete hydrolysis of binary compounds:

Al 2 S 3 + 6H 2 O = 2Al(OH) 3 ↓ + 3 H2S

Sodium sulfide Na 2 S. Oxygen-free salt. White, very hygroscopic. Melts without decomposition, thermally stable. It is highly soluble in water, hydrolyzes at the anion, and creates a highly alkaline environment in solution. When exposed to air, the solution becomes cloudy (colloidal sulfur) and turns yellow (polysulfide color). Typical reducer. Adds sulfur. Enters into ion exchange reactions.

Qualitative reactions on the S 2‑ ion – precipitation of differently colored metal sulfides, of which MnS, FeS, ZnS decompose into HCl (diluted).

It is used in the production of sulfur dyes and cellulose, for removing hair from hides when tanning leather, as a reagent in analytical chemistry.

Equations of the most important reactions:

Na 2 S + 2НCl (diluted) = 2NaCl + H 2 S

Na 2 S + 3H 2 SO 4 (conc.) = SO 2 + S↓ + 2H 2 O + 2NaHSO 4 (up to 50 °C)

Na 2 S + 4HNO 3 (conc.) = 2NO + S↓ + 2H 2 O + 2NaNO 3 (60 °C)

Na 2 S + H 2 S (saturated) = 2NaHS

Na 2 S (t) + 2O 2 = Na 2 SO 4 (above 400 °C)

Na 2 S + 4H 2 O 2 (conc.) = Na 2 SO 4 + 4H 2 O

S 2‑ + M 2+ = MnS (tel.)↓; FeS (black)↓; ZnS (white)↓

S 2‑ + 2Ag + = Ag 2 S (black)↓

S 2‑ + M 2+ = СdS (yellow)↓; PbS, CuS, HgS (black)↓

3S 2‑ + 2Bi 3+ = Bi 2 S 3 (cor. – black)↓

3S 2‑ + 6H 2 O + 2M 3+ = 3H 2 S + 2M(OH) 3 ↓ (M = Al, Cr)

Receipt V industry– calcination of the mineral mirabilite Na 2 SO 4 10H 2 O in the presence of reducing agents:

Na 2 SO 4 + 4H 2 = Na 2 S + 4H 2 O (500 °C, cat. Fe 2 O 3)

Na 2 SO 4 + 4С (coke) = Na 2 S + 4СО (800–1000 °C)

Na 2 SO 4 + 4СО = Na 2 S + 4СО 2 (600–700 °C)

Aluminum sulfide Al 2 S 3. Oxygen-free salt. White, the Al–S bond is predominantly covalent. Melts without decomposition under excess pressure N 2, easily sublimes. Oxidizes in air when heated. It is completely hydrolyzed by water and does not precipitate from solution. Decomposes with strong acids. Used as a solid source of pure hydrogen sulfide. Equations of the most important reactions:

Al 2 S 3 + 6H 2 O = 2Al(OH) 3 ↓ + 3H 2 S (pure)

Al 2 S 3 + 6HCl (diluted) = 2AlCl 3 + 3H 2 S

Al 2 S 3 + 24HNO 3 (conc.) = Al 2 (SO 4) 3 + 24NO 2 + 12H 2 O (100 °C)

2Al 2 S 3 + 9O 2 (air) = 2Al 2 O 3 + 6SO 2 (700–800 °C)

Receipt: interaction of aluminum with molten sulfur in the absence of oxygen and moisture:

2Al + 3S = AL 2 S 3(150–200 °C)

Iron (II) sulfide FeS. Oxygen-free salt. Black-gray with a green tint, refractory, decomposes when heated in a vacuum. When wet, it is sensitive to air oxygen. Insoluble in water. Does not precipitate when solutions of iron(II) salts are saturated with hydrogen sulfide. Decomposes with acids. It is used as a raw material in the production of cast iron, a solid source of hydrogen sulfide.

The iron(III) compound Fe 2 S 3 is not known (not obtained).

Equations of the most important reactions:

Receipt:

Fe + S = FeS(600 °C)

Fe 2 O 3 + H 2 + 2H 2 S = 9 FeS+ 3H 2 O (700‑1000 °C)

FeCl 2 + 2NH 4 HS (g) = FeS↓ + 2NH 4 Cl + H 2 S

Iron disulfide FeS 2. Binary connection. It has the ionic structure Fe 2+ (–S – S–) 2‑ . Dark yellow, thermally stable, decomposes when heated. Insoluble in water, does not react with dilute acids and alkalis. Decomposes by oxidizing acids and is fired in air. It is used as a raw material in the production of cast iron, sulfur and sulfuric acid, and a catalyst in organic synthesis. Ore minerals found in nature pyrite And Marcasite.

Equations of the most important reactions:

FeS 2 = FeS + S (above 1170 °C, vacuum)

2FeS 2 + 14H 2 SO 4 (conc., horizontal) = Fe 2 (SO 4) 3 + 15SO 2 + 14H 2 O

FeS 2 + 18HNO 3 (conc.) = Fe(NO 3) 3 + 2H 2 SO 4 + 15NO 2 + 7H 2 O

4FeS 2 + 11O 2 (air) = 8SO 2 + 2Fe 2 O 3 (800 °C, roasting)

Ammonium hydrosulfide NH 4 HS. An oxygen-free acidic salt. White, melts under excess pressure. Very volatile, thermally unstable. It oxidizes in air. It is highly soluble in water, hydrolyzes into the cation and anion (predominates), creates an alkaline environment. The solution turns yellow in air. Decomposes with acids and adds sulfur in a saturated solution. It is not neutralized by alkalis, the average salt (NH 4) 2 S does not exist in solution (for the conditions for obtaining the average salt, see the “H 2 S” section). It is used as a component of photographic developers, as an analytical reagent (sulfide precipitator).

Equations of the most important reactions:

NH 4 HS = NH 3 + H 2 S (above 20 °C)

NH 4 HS + HCl (diluted) = NH 4 Cl + H 2 S

NH 4 HS + 3HNO 3 (conc.) = S↓ + 2NO 2 + NH 4 NO 3 + 2H 2 O

2NH 4 HS (saturated H 2 S) + 2CuSO 4 = (NH 4) 2 SO 4 + H 2 SO 4 + 2CuS↓

Receipt: saturation of a concentrated solution of NH 3 with hydrogen sulfide:

NH 3 H 2 O (conc.) + H 2 S (g) = NH 4 HS+ H 2 O

In analytical chemistry, a solution containing equal amounts of NH 4 HS and NH 3 H 2 O is conventionally considered a solution of (NH 4) 2 S and the formula of the average salt is used in writing the reaction equations, although ammonium sulfide is completely hydrolyzed in water to NH 4 HS and NH 3H2O.

Sulfur dioxide. Sulfites

Sulfur dioxide SO2. Acidic oxide. Colorless gas with a pungent odor. The molecule has the structure of an incomplete triangle [: S(O) 2 ] (sp 2 - hybridization), contains σ, π bonds S=O. Easily liquefied, thermally stable. Highly soluble in water (~40 l/1 l H 2 O at 20 °C). Forms a polyhydrate with the properties of a weak acid; dissociation products are HSO 3 - and SO 3 2 - ions. The HSO 3 ion has two tautomeric forms - symmetrical(non-acidic) with a tetrahedral structure (sp 3 -hybridization), which predominates in the mixture, and asymmetrical(acidic) with the structure of an incomplete tetrahedron [: S(O) 2 (OH)] (sp 3 -hybridization). The SO 3 2‑ ion is also tetrahedral [: S(O) 3 ].

Reacts with alkalis, ammonia hydrate. A typical reducing agent, weak oxidizing agent.

Qualitative reaction– discoloration of yellow-brown “iodine water”. Intermediate product in the production of sulfites and sulfuric acid.

It is used for bleaching wool, silk and straw, canning and storing fruits, as a disinfectant, antioxidant, and refrigerant. Poisonous.

The compound of composition H 2 SO 3 (sulfurous acid) is unknown (does not exist).

Equations of the most important reactions:

Solubility in water and acidic properties:

Receipt: in industry - combustion of sulfur in air enriched with oxygen, and, to a lesser extent, roasting of sulfide ores (SO 2 - associated gas when roasting pyrite):

S + O 2 = SO 2(280–360 °C)

4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8 SO 2(800 °C, firing)

in the laboratory - displacement of sulfites with sulfuric acid:

BaSO 3 (t) + H 2 SO 4 (conc.) = BaSO 4 ↓ + SO 2 + H 2 O

Sodium sulfite Na 2 SO 3. Oxosol. White. When heated in air, it decomposes without melting and melts under excess pressure of argon. When wet and in solution, it is sensitive to atmospheric oxygen. It is highly soluble in water and hydrolyzes at the anion. Decomposes with acids. Typical reducer.

Qualitative reaction on the SO 3 2‑ ion - the formation of a white precipitate of barium sulfite, which is transferred into solution with strong acids (HCl, HNO 3).

It is used as a reagent in analytical chemistry, a component of photographic solutions, and a chlorine neutralizer for bleaching fabrics.

Equations of the most important reactions:

Receipt:

Na 2 CO 3 (conc.) + SO 2 = Na2SO3+CO2

Sulfuric acid. Sulfates

Sulfuric acid H 2 SO 4. Oxoacid. Colorless liquid, very viscous (oily), very hygroscopic. The molecule has a distorted tetrahedral structure (sp 3 -hybridization), contains covalent σ-bonds S – OH and σπ-bonds S=O. The SO 4 2‑ ion has a regular tetrahedral structure. It has a wide temperature range of the liquid state (~300 degrees). Partially decomposes when heated above 296 °C. It is distilled in the form of an azeotropic mixture with water (mass fraction of acid 98.3%, boiling point 296–340 °C), and with stronger heating it decomposes completely. Unlimitedly miscible with water (with strong exo‑effect). Strong acid in solution, neutralized by alkalis and ammonia hydrate. Converts metals into sulfates (with an excess of concentrated acid under normal conditions, soluble hydrosulfates are formed), but the metals Be, Bi, Co, Fe, Mg and Nb are passivated in concentrated acid and do not react with it. Reacts with basic oxides and hydroxides, decomposes salts of weak acids. A weak oxidizing agent in a dilute solution (due to H I), a strong oxidizing agent in a concentrated solution (due to S VI). It dissolves SO 3 well and reacts with it (a heavy oily liquid is formed - oleum, contains H 2 S 2 O 7).

Qualitative reaction on the SO 4 2‑ ion – precipitation of white barium sulfate BaSO 4 (the precipitate is not transferred into solution by hydrochloric and nitric acids, unlike the white precipitate BaSO 3).

It is used in the production of sulfates and other sulfur compounds, mineral fertilizers, explosives, dyes and drugs, in organic synthesis, for the “opening” (the first stage of processing) of industrially important ores and minerals, in the purification of petroleum products, in the electrolysis of water, as an electrolyte for lead batteries . Toxic, causes skin burns. Equations of the most important reactions:

Receipt V industry:

a) synthesis of SO 2 from sulfur, sulfide ores, hydrogen sulfide and sulfate ores:

S + O 2 (air) = SO 2(280–360 °C)

4FeS 2 + 11O 2 (air) = 8 SO 2+ 2Fe 2 O 3 (800 °C, firing)

2H 2 S + 3O 2 (g) = 2 SO 2+ 2H 2 O (250–300 °C)

CaSO 4 + C (coke) = CaO + SO 2+ CO (1300–1500 °C)

b) conversion of SO 2 to SO 3 in a contact apparatus:

c) synthesis of concentrated and anhydrous sulfuric acid:

H 2 O (dil. H 2 SO 4) + SO 3 = H2SO4(conc., anhydrous)

(the absorption of SO 3 with pure water to produce H 2 SO 4 is not carried out due to the strong heating of the mixture and the reverse decomposition of H 2 SO 4, see above);

d) synthesis oleum– a mixture of anhydrous H 2 SO 4, disulfuric acid H 2 S 2 O 7 and excess SO 3. Dissolved SO 3 guarantees the anhydrity of oleum (when water enters, H 2 SO 4 is immediately formed), which allows it to be safely transported in steel tanks.

Sodium sulfate Na 2 SO 4. Oxosol. White, hygroscopic. Melts and boils without decomposition. Forms crystalline hydrate (mineral mirabilite), easily losing water; technical name Glauber's salt. It is highly soluble in water and does not hydrolyze. Reacts with H 2 SO 4 (conc.), SO 3 . It is reduced by hydrogen and coke when heated. Enters into ion exchange reactions.

It is used in the production of glass, cellulose and mineral paints, as a medicine. Contained in the brine of salt lakes, in particular in the Kara-Bogaz-Gol Bay of the Caspian Sea.

Equations of the most important reactions:

Potassium hydrogen sulfate KHSO 4. Acid oxo salt. White, hygroscopic, but does not form crystalline hydrates. When heated, it melts and decomposes. It is highly soluble in water; the anion undergoes dissociation in solution; the solution environment is strongly acidic. Neutralized by alkalis.

It is used as a component of fluxes in metallurgy, an integral part of mineral fertilizers.

Equations of the most important reactions:

2KHSO 4 = K 2 SO 4 + H 2 SO 4 (up to 240 °C)

2KHSO 4 = K 2 S 2 O 7 + H 2 O (320–340 °C)

KHSO 4 (dil.) + KOH (conc.) = K 2 SO 4 + H 2 O KHSO 4 + KCl = K 2 SO 4 + HCl (450–700 °C)

6KHSO 4 + M 2 O 3 = 2KM(SO 4) 2 + 2K 2 SO 4 + 3H 2 O (350–500 °C, M = Al, Cr)

Receipt: treatment of potassium sulfate with concentrated (more than 6O%) sulfuric acid in the cold:

K 2 SO 4 + H 2 SO 4 (conc.) = 2 KHSO 4

Calcium sulfate CaSO 4. Oxosol. White, very hygroscopic, refractory, decomposes when heated. Natural CaSO 4 occurs as a very common mineral gypsum CaSO 4 2H 2 O. At 130 °C, gypsum loses some of the water and turns into burnt (plaster) gypsum 2CaSO 4 H 2 O (technical name alabaster). Completely dehydrated (200 °C) gypsum corresponds to the mineral anhydrite CaSO4. Slightly soluble in water (0.206 g/100 g H 2 O at 20 °C), solubility decreases when heated. Reacts with H 2 SO 4 (conc.). Restored by coke during fusion. Determines most of the “permanent” hardness of fresh water (see 9.2 for details).

Equations of the most important reactions: 100–128 °C

It is used as a raw material in the production of SO 2, H 2 SO 4 and (NH 4) 2 SO 4, as a flux in metallurgy, and as a paper filler. A binder mortar made from burnt gypsum “sets” faster than a mixture based on Ca(OH) 2 . Hardening is ensured by the binding of water, the formation of gypsum in the form of a stone mass. Burnt gypsum is used to make plaster casts, architectural and decorative forms and products, partition slabs and panels, and stone floors.

Aluminum-potassium sulfate KAl(SO 4) 2. Double oxo salt. White, hygroscopic. Decomposes when heated strongly. Forms crystalline hydrate - potassium alum. Moderately soluble in water, hydrolyzes with aluminum cation. Reacts with alkalis, ammonia hydrate.

It is used as a mordant for dyeing fabrics, a leather tanning agent, a coagulant for purifying fresh water, a component of compositions for sizing paper, and an external hemostatic agent in medicine and cosmetology. It is formed by the joint crystallization of aluminum and potassium sulfates.

Equations of the most important reactions:

Chromium(III) sulfate - potassium KCr(SO 4) 2. Double oxo salt. Red (hydrate dark purple, technical name chromium-potassium alum). When heated, it decomposes without melting. It is highly soluble in water (the gray-blue color of the solution corresponds to aqua complex 3+), hydrolyzes at the chromium(III) cation. Reacts with alkalis, ammonia hydrate. Weak oxidizing and reducing agent. Enters into ion exchange reactions.

Qualitative reactions on the Cr 3+ ion – reduction to Cr 2+ or oxidation to yellow CrO 4 2‑.

It is used as a leather tanning agent, a mordant for dyeing fabrics, and a reagent in photography. It is formed by the joint crystallization of chromium(III) and potassium sulfates. Equations of the most important reactions:

Manganese (II) sulfate MnSO 4 . Oxosol. White, melts and decomposes when heated. Crystalline hydrate MnSO 4 5H 2 O – red-pink, technical name manganese sulfate. It is highly soluble in water; the light pink (almost colorless) color of the solution corresponds to aquacomplex 2+; hydrolyzes at the cation. Reacts with alkalis, ammonia hydrate. Weak reducing agent, reacts with typical (strong) oxidizing agents.

Qualitative reactions on the Mn 2+ ion – commutation with the MnO 4 ion and the disappearance of the violet color of the latter, oxidation of Mn 2+ to MnO 4 and the appearance of a violet color.

It is used for the production of Mn, MnO 2 and other manganese compounds, as a microfertilizer and analytical reagent.

Equations of the most important reactions:

Receipt:

2MnO 2 + 2H 2 SO 4 (conc.) = 2 MnSO4+ O 2 + 2H 2 O (100 °C)

Iron (II) sulfate FeSO 4 . Oxosol. White (light green hydrate, technical name inkstone), hygroscopic. Decomposes when heated. It is highly soluble in water and is slightly hydrolyzed by the cation. It is quickly oxidized in solution by atmospheric oxygen (the solution turns yellow and becomes cloudy). Reacts with oxidizing acids, alkalis, and ammonia hydrate. Typical reducer.

It is used as a component of mineral paints, electrolytes in electroplating, a wood preservative, a fungicide, and a medicine against anemia. In the laboratory it is often taken in the form of a double salt Fe(NH 4) 2 (SO 4) 2 6H 2 O ( Mohr's salt), more resistant to air.

Equations of the most important reactions:

Receipt:

Fe + H 2 SO 4 (diluted) = FeSO4+H2

FeCO 3 + H 2 SO 4 (diluted) = FeSO4+ CO 2 + H 2 O

7.4. Non-metals VA‑group

Nitrogen. Ammonia

Nitrogen– element of the 2nd period and VA group of the Periodic system, serial number 7. Electronic formula of the atom [ 2 He]2s 2 2p 3, characteristic oxidation states 0, ‑III, +III and +V, less often +II, +IV and etc.; the N v state is considered relatively stable.

Scale of nitrogen oxidation states:

Nitrogen has a high electronegativity (3.07), third after F and O. It exhibits typical non-metallic (acidic) properties. Forms various oxygen-containing acids, salts and binary compounds, as well as the ammonium cation NH 4 + and its salts.

In nature - seventeenth by chemical abundance element (ninth among non-metals). A vital element for all organisms.

Nitrogen N 2. Simple substance. It consists of non-polar molecules with a very stable σππ-bond N ≡ N, this explains the chemical inertness of nitrogen under normal conditions. A colorless, tasteless and odorless gas that condenses into a colorless liquid (unlike O2).

Main component of air: 78.09% by volume, 75.52% by mass. Nitrogen boils away from liquid air before oxygen O2. Slightly soluble in water (15.4 ml/1 l H 2 O at 20 ° C), the solubility of nitrogen is less than that of oxygen.

At room temperature, N2 reacts only with lithium (in a humid atmosphere), forming lithium nitride Li3N; nitrides of other elements are synthesized with strong heating:

N 2 + 3Mg = Mg 3 N 2 (800 °C)

In an electrical discharge, N2 reacts with fluorine and, to a very small extent, with oxygen:

The reversible reaction to produce ammonia occurs at 500 °C, under pressure up to 350 atm and always in the presence of a catalyst (Fe/F 2 O 3 /FeO, in the laboratory Pt):

According to Le Chatelier's principle, an increase in ammonia yield should occur with increasing pressure and decreasing temperature. However, the reaction rate at low temperatures is very low, so the process is carried out at 450–500 °C, achieving a 15% ammonia yield. Unreacted N 2 and H 2 are returned to the reactor and thereby increase the degree of reaction.

Nitrogen is chemically passive in relation to acids and alkalis and does not support combustion.

Receipt V industry– fractional distillation of liquid air or removal of oxygen from air by chemical means, for example, by the reaction 2C (coke) + O 2 = 2CO when heated. In these cases, nitrogen is obtained, which also contains impurities of noble gases (mainly argon).

IN laboratories small amounts of chemically pure nitrogen can be obtained by the commutation reaction with moderate heating:

N ‑III H 4 N III O 2(t) = N 2 0 + 2H 2 O (60–70 °C)

NH 4 Cl (p) + KNO 2 (p) = N 2 0 + KCl + 2H 2 O (100 °C)

It is used for the synthesis of ammonia, nitric acid and other nitrogen-containing products, as an inert medium for chemical and metallurgical processes and storage of flammable substances.

Ammonia NH3. Binary compound, the oxidation state of nitrogen is – III. Colorless gas with a sharp characteristic odor. The molecule has the structure of an incomplete tetrahedron [: N(H) 3)] (sp 3 -hybridization). The presence of a donor pair of electrons on the sp 3 -hybrid orbital of nitrogen in the NH 3 molecule determines the characteristic reaction of addition of a hydrogen cation, which results in the formation of a cation ammonium NH4+. It liquefies under excess pressure at room temperature. In the liquid state, it is associated through hydrogen bonds. Thermally unstable. Highly soluble in water (more than 700 l/1 l H 2 O at 20 °C); the proportion in the saturated solution is = 34% by mass and = 99% by volume, pH = 11.8.

Very reactive, prone to addition reactions. Cr reacts in oxygen, reacts with acids. It exhibits reducing (due to N‑III) and oxidizing (due to H I) properties. It is dried only with calcium oxide.

Qualitative reactions– formation of white “smoke” upon contact with gaseous HCl, blackening of a piece of paper moistened with a solution of Hg 2 (NO 3) 2.

An intermediate product in the synthesis of HNO 3 and ammonium salts. Used in the production of soda, nitrogen fertilizers, dyes, explosives; liquid ammonia is a refrigerant. Poisonous.

Equations of the most important reactions:

Receipt: V laboratories– displacement of ammonia from ammonium salts when heated with soda lime (NaOH + CaO):

or boiling an aqueous solution of ammonia and then drying the gas.

IN industry ammonia is synthesized from nitrogen (see) with hydrogen. Produced by industry either in liquefied form or in the form of a concentrated aqueous solution under the technical name ammonia water.

Ammonia hydrate NH 3 H 2 O. Intermolecular connection. White, in the crystal lattice - molecules NH 3 and H 2 O, connected by a weak hydrogen bond H 3 N ... HON. Present in an aqueous solution of ammonia, a weak base (dissociation products - NH 4 ‑ cation and OH ‑ anion). The ammonium cation has a regular tetrahedral structure (sp 3 hybridization). Thermally unstable, completely decomposes when the solution is boiled. Neutralized by strong acids. Shows reducing properties (due to N III) in a concentrated solution. Enters into ion exchange and complexation reactions.

Qualitative reaction– formation of white “smoke” upon contact with gaseous HCl.

It is used to create a slightly alkaline environment in solution during the precipitation of amphoteric hydroxides.

A 1M ammonia solution contains mainly NH 3 H 2 O hydrate and only 0.4% NH 4 + and OH - ions (due to hydrate dissociation); Thus, the ionic “ammonium hydroxide NH 4 OH” is practically not contained in the solution, and there is no such compound in the solid hydrate. Equations of the most important reactions:

NH 3 H 2 O (conc.) = NH 3 + H 2 O (boiling with NaOH)

NH 3 H 2 O + HCl (diluted) = NH 4 Cl + H 2 O

3(NH 3 H 2 O) (conc.) + CrCl 3 = Cr(OH) 3 ↓ + 3NH 4 Cl

8(NH 3 H 2 O) (conc.) + ZBr 2 (p) = N 2 + 6NH 4 Br + 8H 2 O (40–50 °C)

2(NH 3 H 2 O) (conc.) + 2KMnO 4 = N 2 + 2MnO 2 ↓ + 4H 2 O + 2KOH

4(NH 3 H 2 O) (conc.) + Ag2O= 2OH + 3H2O

4(NH 3 H 2 O) (conc.) + Cu(OH) 2 + (OH) 2 + 4H 2 O

6(NH 3 H 2 O) (conc.) + NiCl 2 = Cl 2 + 6H 2 O

A dilute ammonia solution (3–10%) is often called ammonia(the name was invented by alchemists), and the concentrated solution (18.5–25%) - ammonia water(produced by industry).

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