Obtaining hydrogen sulfide. Production of sulfur dioxide by burning sulfur, hydrogen sulfide and other types of raw materials Hydrogen sulfide sulfur dioxide

Almurzinova Zavrish Bisembaevna , teacher of biology and chemistry MBOU “State Farm Basic Secondary School of Adamovsky District, Orenburg Region.

Subject - chemistry, grade - 9.

Educational complex: “Inorganic chemistry”, authors: G.E. Rudzitis, F.G. Feldman, Moscow, “Enlightenment”, 2014.

Level of training – basic.

Subject : “Hydrogen sulfide. Sulfides. Sulphur dioxide. Sulfurous acid and its salts." Number of hours on the topic – 1.

Lesson No. 4 in the lesson system on the topic« Oxygen and sulfur ».

Target : Based on knowledge of the structure of hydrogen sulfide and sulfur oxides, consider their properties and production, introduce students to methods for recognizing sulfides and sulfites.

Tasks:

1. Educational – study the structural features and properties of sulfur compounds (II) And(IV); become familiar with qualitative reactions to sulfide and sulfite ions.

2. Developmental – develop students’ skills in conducting experiments, observing results, analyzing and drawing conclusions.

3. Educational – developing interest in what is being studied, instilling skills in relating to nature.

Planned results : be able to describe the physical and chemical properties of hydrogen sulfide, hydrogen sulfide acid and its salts; know methods for producing sulfur dioxide and sulfurous acid, explain the properties of sulfur compounds(II) and (IV) based on ideas about redox processes; have an idea of the effect of sulfur dioxide on the occurrence of acid rain.

Equipment : On the demonstration table: sulfur, sodium sulfide, iron sulfide, litmus solution, sulfuric acid solution, lead nitrate solution, chlorine in a cylinder closed with a stopper, a device for producing hydrogen sulfide and testing its properties, sulfur oxide (VI), oxygen gas meter, 500 ml glass, spoon for burning substances.

During the classes :

Organizing time .

We conduct a conversation on repeating the properties of sulfur:

1) what explains the presence of several allotropic modifications of sulfur?

2) what happens to the molecules: A) when vaporous sulfur is cooled. B) during long-term storage of plastic sulfur, c) when crystals precipitate from a solution of sulfur in organic solvents, for example in toluene?

3) what is the flotation method of purifying sulfur from impurities, for example from river sand, based on?

We call two students: 1) draw diagrams of molecules of various allotropic modifications of sulfur and talk about their physical properties. 2) compose reaction equations characterizing the properties of oxygen and consider them from the point of view of oxidation-reduction.

The rest of the students solve the problem: what is the mass of zinc sulfide formed during the reaction of a zinc compound with sulfur, taken with an amount of substance of 2.5 mol?

Together with the students, we formulate the lesson objective : get acquainted with the properties of sulfur compounds with oxidation states -2 and +4.

New topic : Students name compounds known to them in which sulfur exhibits these oxidation states. Chemical, electronic and structural formulas of hydrogen sulfide and sulfur oxide (IV), sulfurous acid.

How can you get hydrogen sulfide? Students write down the equation for the reaction of sulfur with hydrogen and explain it from the point of view of oxidation-reduction. Then another method for producing hydrogen sulfide is considered: the exchange reaction of acids with metal sulfides. Let's compare this method with methods for producing hydrogen halides. We note that the degree of sulfur oxidation in exchange reactions does not change.

What properties does hydrogen sulfide have? In a conversation, we find out the physical properties and note the physiological effect. We determine the chemical properties by experimenting with the combustion of hydrogen sulfide in air under various conditions. What can be formed as reaction products? We consider reactions from the point of view of oxidation-reduction:

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

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

We draw students' attention to the fact that with complete combustion, more complete oxidation occurs (S -2 - 6 e - = S +4 ) than in the second case (S -2 - 2 e - = S 0 ).

We discuss how the process will go if chlorine is used as an oxidizing agent. We demonstrate the experience of mixing gases in two cylinders, the top of which is pre-filled with chlorine, the bottom with hydrogen sulfide. Chlorine becomes discolored and hydrogen chloride is formed. Sulfur settles on the walls of the cylinder. After this, we consider the essence of the decomposition reaction of hydrogen sulfide and lead students to the conclusion about the acidic nature of hydrogen sulfide, confirming it with experience with litmus. Then we carry out a qualitative reaction to the sulfide ion and compose the reaction equation:

Na 2 S+Pb(NO 3 ) 2 =2NaNO 3 +PbS ↓

Together with the students, we formulate the conclusion: hydrogen sulfide is only a reducing agent in redox reactions, it is acidic in nature, and its solution in water is an acid.

S 0 →S -2 ; S -2 →S 0 ; S 0 →S +4 ; S -2 →S +4 ; S 0 →H 2 S -2 → S +4 ABOUT 2.

We lead students to the conclusion that there is a genetic connection between sulfur compounds and begin a conversation about the compoundsS +4 . We demonstrate experiments: 1) obtaining sulfur oxide (IV), 2) discoloration of the fuchsin solution, 3) dissolution of sulfur oxide (IV) in water, 4) acid detection. We compose reaction equations for the experiments performed and analyze the essence of the reactions:

2SABOUT 2 + ABOUT 2 =2 SABOUT 3 ; SABOUT 2 +2H 2 S=3S+2H 2 ABOUT.

Sulfurous acid is an unstable compound, easily decomposes into sulfur oxide (IV) and water, therefore it exists only in aqueous solutions. This acid is of medium strength. It forms two rows of salts: the middle ones are sulfites (SABOUT 3 -2 ), acidic – hydrosulfites (H.S.ABOUT 3 -1 ).

We demonstrate experience: qualitative determination of sulfites, interaction of sulfites with a strong acid, which releases gasSABOUT 2 pungent odor:

TO 2 SABOUT 3 + N 2 SABOUT 4 → K 2 SABOUT 4 + N 2 O +SABOUT 2

Consolidation. Work on two options to draw up application schemes: 1 option for hydrogen sulfide, the second option for sulfur oxide (IV)

Reflection . Let's summarize the work:

What connections did we talk about today?

What properties do sulfur compounds exhibit?II) And (IV).

Name the areas of application of these compounds

VII. Homework: §11,12, exercises 3-5 (p.34)

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§ 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 ultraviolet rays of the Sun, 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 ozonator. 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 does potassium iodide solution need to 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. Reductive 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 the 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 produced 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 releasing 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?

Chemical properties

Physical properties

Under normal conditions, hydrogen sulfide is a colorless gas with a strong, characteristic odor of rotten eggs. T pl = -86 °C, T kip = -60 °C, poorly soluble in water, at 20 °C 2.58 ml of H 2 S dissolves in 100 g of water. Very toxic, if inhaled it causes paralysis, which can be fatal. In nature, it is released as part of volcanic gases and is formed during the decay of plant and animal organisms. It is highly soluble in water; when dissolved, it forms weak hydrosulfide acid.

In an aqueous solution, hydrogen sulfide has the properties of a weak dibasic acid:

H 2 S = HS - + H + ;

HS - = S 2- + H + .

Hydrogen sulfide burns in the air blue flame. With limited air access, free sulfur is formed:

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

With excess air access, combustion of hydrogen sulfide leads to the formation of sulfur oxide (IV):

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

Hydrogen sulfide has reducing properties. Depending on conditions, hydrogen sulfide can be oxidized in aqueous solution to sulfur, sulfur dioxide and sulfuric acid.

For example, it decolorizes bromine water:

H 2 S + Br 2 = 2HBr + S.

interacts with chlorine water:

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

A stream of hydrogen sulfide can be ignited using lead dioxide, since the reaction is accompanied by a large release of heat:

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

Interaction of hydrogen sulfide with sulfur dioxide used to obtain sulfur from waste gases of metallurgical and sulfuric acid production:

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

The formation of native sulfur during volcanic processes is associated with this process.

When sulfur dioxide and hydrogen sulfide are simultaneously passed through an alkali solution, thiosulfate is formed:

4SO 2 + 2H 2 S + 6NaOH = 3Na 2 S 2 O 3 + 5H 2 O.

Reaction of dilute hydrochloric acid with iron (II) sulfide

FeS + 2HCl = FeCl 2 + H 2 S

Reaction of aluminum sulfide with cold water

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

Direct synthesis from elements occurs when hydrogen is passed over molten sulfur:

H 2 + S = H 2 S.

Heating a mixture of paraffin and sulfur.

1.9. Hydrogen sulfide acid and its salts

Hydrogen sulfide acid has all the properties of weak acids. It reacts with metals, metal oxides, and bases.

As a dibasic acid, it forms two types of salts - sulfides and hydrosulfides . Hydrosulfides are highly soluble in water, sulfides of alkali and alkaline earth metals as well, and sulfides of heavy metals are practically insoluble.

Sulfides of alkali and alkaline earth metals are not colored, the rest have a characteristic color, for example, sulfides of copper (II), nickel and lead - black, cadmium, indium, tin - yellow, antimony - orange.

Ionic alkali metal sulfides M 2 S have a fluorite-type structure, where each sulfur atom is surrounded by a cube of 8 metal atoms and each metal atom is surrounded by a tetrahedron of 4 sulfur atoms. MS-type sulfides are characteristic of alkaline earth metals and have a sodium chloride-type structure, where each metal and sulfur atom is surrounded by an octahedron of atoms of a different type. When the covalent nature of the metal–sulfur bond is strengthened, structures with lower coordination numbers are realized.

Sulfides of non-ferrous metals are found in nature as minerals and ores and serve as raw materials for the production of metals.

Chemistry tutor

Continuation. See in No. 22/2005; 1, 2, 3, 5, 6, 8, 9, 11, 13, 15, 16, 18, 22/2006;
3, 4, 7, 10, 11, 21/2007;
2, 7, 11, 18, 19, 21/2008;
1, 3, 10/2009

LESSON 30

10th grade (first year of study)

Sulfur and its compounds

1. Position in the table of D.I. Mendeleev, structure of the atom.

2. Origin of the name.

3. Physical properties.

4. Chemical properties.

5. Being in nature.

6. Basic methods of obtaining.

7. The most important sulfur compounds (hydrogen sulfide, hydrosulfide acid and its salts; sulfur dioxide, sulfurous acid and its salts; sulfur trioxide, sulfuric acid and its salts).

In the periodic table, sulfur is in the main subgroup of group VI (chalcogen subgroup). Electronic formula of sulfur 1 s 2 2s 2 p 6 3s 2 p 4, this R-element. Depending on its state, sulfur can exhibit valency II, IV or VI:

S: 1 s 2 2s 2 2p 6 3s 2 3p 4 3d 0 (valency II),

S*: 1 s 2 2s 2 2p 6 3s 2 3p 3 3d 1 (valence IV),

S**: 1 s 2 2s 2 2p 6 3s 1 3p 3 3d 2 (valency VI).

The characteristic oxidation states of sulfur are –2, +2, +4, +6 (in disulfides containing a bridged –S–S– bond (for example, FeS 2), the oxidation state of sulfur is –1); in compounds it is part of anions, with more electronegative elements – part of cations, for example:

Sulfur – an element with high electronegativity, exhibits non-metallic (acidic) properties. It has four stable isotopes with mass numbers 32, 33, 34 and 36. Natural sulfur is 95% composed of the 32 S isotope.

The Russian name for sulfur comes from the Sanskrit word cira– light yellow, the color of natural sulfur. Latin name sulfur translated as "flammable powder". 1

PHYSICAL STRUCTURES

Sulfur forms three allotropic modifications: rhombic(-sulfur), monoclinic(-sulfur) and plastic, or rubbery. Orthorhombic sulfur is most stable under normal conditions, and monoclinic sulfur is stable above 95.5 °C. Both of these allotropic modifications have a molecular crystal lattice built from molecules of the composition S 8 located in space in the form of a crown; atoms are connected by single covalent bonds. The difference between rhombic and monoclinic sulfur is that in the crystal lattice the molecules are packed differently.

If rhombic or monoclinic sulfur is heated to its boiling point (444.6 °C) and the resulting liquid is poured into cold water, plastic sulfur is formed, with properties resembling rubber. Plastic sulfur consists of long zigzag chains. This allotropic modification is unstable and spontaneously transforms into one of the crystalline forms.

Rhombic sulfur is a yellow crystalline solid; does not dissolve in water (and is not wetted), but is highly soluble in many organic solvents (carbon disulfide, benzene, etc.). Sulfur has very poor electrical and thermal conductivity. The melting point of orthorhombic sulfur is +112.8 °C; at a temperature of 95.5 °C, orthorhombic sulfur becomes monoclinic:

Chemical properties

In terms of its chemical properties, sulfur is a typical active non-metal. In reactions it can be both an oxidizing agent and a reducing agent.

Metals (+):

2Na + S = Na 2 S,

2Al + 3S Al 2 S 3,

Non-metals (+/–)*:

2P + 3S P 2 S 3 ,

S + Cl 2 = SCl 2,

S + 3F 2 = SF 6,

S + N 2 reaction does not occur.

H 2 O (–). sulfur is not wetted by water.

Basic oxides (–).

Acidic oxides (–).

Bases (+/–):

S + Cu(OH) 2 reaction does not occur.

Acids (not oxidizing agents) (–).

Oxidizing acids (+):

S + 2H 2 SO 4 (conc.) = 3SO 2 + 2H 2 O,

S + 2HNO 3 (diluted) = H 2 SO 4 + 2NO,

S + 6HNO 3 (conc.) = H 2 SO 4 + 6NO 2 + 2H 2 O.

In nature, sulfur occurs both in the native state and in the form of compounds, the most important of which are pyrite, also known as iron or sulfur pyrite (FeS 2), zinc blende (ZnS), lead luster (PbS ), gypsum (CaSO 4 2H 2 O), Glauber's salt (Na 2 SO 4 10H 2 O), bitter salt (MgSO 4 7H 2 O). In addition, sulfur is part of coal, oil, as well as various living organisms (as part of amino acids). In the human body, sulfur is concentrated in the hair.

In laboratory conditions, sulfur can be obtained using redox reactions (ORR), for example:

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

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

IMPORTANT SULPHUR COMPOUNDS

Hydrogen sulfide (H 2 S) is a colorless gas with a suffocating, unpleasant odor of rotten eggs, poisonous (combines with hemoglobin in the blood, forming iron sulfide). Heavier than air, slightly soluble in water (2.5 volumes of hydrogen sulfide in 1 volume of water). The bonds in the molecule are polar covalent, sp 3-hybridization, the molecule has an angular structure:

Chemically, hydrogen sulfide is quite active. It is thermally unstable; burns easily in an oxygen atmosphere or in air; easily oxidized by halogens, sulfur dioxide or iron(III) chloride; when heated, it interacts with some metals and their oxides, forming sulfides:

2H 2 S + O 2 2S + 2H 2 O,

2H 2 S + 3O 2 2SO 2 + 2H 2 O,

H 2 S + Br 2 = 2HBr + S,

2H 2 S + SO 2 3S + 2H 2 O,

2FeCl 3 + H 2 S = 2FeCl 2 + S + 2HCl,

H 2 S + Zn ZnS + H 2 ,

H 2 S + CaO CaS + H 2 O.

In laboratory conditions, hydrogen sulfide is obtained by treating iron or zinc sulfides with strong mineral acids or by irreversible hydrolysis of aluminum sulfide:

ZnS + 2HCl = ZnCl 2 + H 2 S,

Al 2 SO 3 + 6HOH 2Al(OH) 3 + 3H 2 S.

Hydrogen sulfide solution in water – hydrogen sulfide water, or hydrosulfide acid . A weak electrolyte, practically does not dissociate in the second stage. How a dibasic acid forms two types of salts − sulfides and hydrosulfides:

for example, Na 2 S – sodium sulfide, NaHS – sodium hydrosulfide.

Hydrogen sulfide acid exhibits all the general properties of acids. In addition, hydrogen sulfide, hydrosulfide acid and its salts exhibit strong reducing ability. For example:

H 2 S + Zn = ZnS + H 2,

H 2 S + CuO = CuS + H 2 O,

Qualitative reaction to sulfide ion is interaction with soluble lead salts; In this case, a black precipitate of lead sulfide precipitates:

Pb 2+ + S 2– -> PbS,

Pb(NO 3) 2 + Na 2 S = PbS + 2NaNO 3.

Sulfur(IV) oxide SO 2 – sulfur dioxide, sulfur dioxide - a colorless gas with a pungent odor, poisonous. Acidic oxide. The bonds in the molecule are polar covalent, sp 2 -hybridization. Heavier than air, highly soluble in water (in one volume of water - up to 80 volumes of SO 2), forms when dissolved sulfurous acid , existing only in solution:

H 2 O + SO 2 H 2 SO 3 .

In terms of acid-base properties, sulfur dioxide exhibits the properties of a typical acid oxide; sulfurous acid also exhibits all the typical properties of acids:

SO 2 + CaO CaSO 3,

H 2 SO 3 + Zn = ZnSO 3 + H 2,

H 2 SO 3 + CaO = CaSO 3 + H 2 O.

In terms of redox properties, sulfur dioxide, sulfurous acid and sulfites can exhibit redox duality (with a predominance of reducing properties). With stronger reducing agents, sulfur(IV) compounds behave as oxidizing agents:

With stronger oxidizing agents they exhibit reducing properties:

IN industry sulfur dioxide is obtained:

When burning sulfur:

Roasting of pyrite and other sulfides:

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

2ZnS + 3O 2 2ZnO + 2SO 2 .

TO laboratory methods receipts include:

The effect of strong acids on sulfites:

Na 2 SO 3 + 2HCl = 2NaCl + SO 2 + H 2 O;

Interaction of concentrated sulfuric acid with heavy metals:

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

Qualitative reactions to sulfite ion– discoloration of “iodine water” or the action of strong mineral acids:

Na 2 SO 3 + I 2 + 2NaOH = 2NaI + Na 2 SO 4 + H 2 O,

Ca 2 SO 3 + 2HCl = CaCl 2 + H 2 O + SO 2.

Sulfur(VI) oxide SO 3 – sulfur trioxide, or sulfuric anhydride , is a colorless liquid, which at temperatures below 17 ° C turns into a white crystalline mass. Poisonous. Exists in the form of polymers (monomer molecules exist only in the gas phase), the bonds in the molecule are polar covalent, sp 2 -hybridization. Hygroscopic, thermally unstable. Reacts with water with a strong exo-effect. Reacts with anhydrous sulfuric acid to form oleum. Formed by the oxidation of sulfur dioxide:

SO 3 + H 2 O = H 2 SO 4 + Q,

n n SO3.

According to its acid-base properties, it is a typical acid oxide:

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

SO 3 + CaO = CaSO 4,

In terms of redox properties, it acts as a strong oxidizing agent, usually being reduced to SO 2 or sulfites:

In its pure form it has no practical value; it is an intermediate product in the production of sulfuric acid.

Sulfuric acid – heavy oily liquid without color and odor. Highly soluble in water (with great exo-effect). Hygroscopic, poisonous, causes severe skin burns. Is a strong electrolyte. Sulfuric acid forms two types of salts: sulfates And hydrosulfates, which exhibit all the general properties of salts. Sulfates of active metals are thermally stable, and sulfates of other metals decompose even with slight heating:

Na 2 SO 4 does not decompose,

ZnSO 4 ZnO + SO 3,

4FeSO 4 2Fe 2 O 3 + 4SO 2 + O 2,

Ag 2 SO 4 2Ag + SO 2 + O 2,

HgSO 4 Hg + SO 2 + O 2.

A solution with a mass fraction of sulfuric acid below 70% is usually considered dilute; above 70% – concentrated; a solution of SO 3 in anhydrous sulfuric acid is called oleum (the concentration of sulfur trioxide in oleum can reach 65%).

Diluted sulfuric acid exhibits all the properties characteristic of strong acids:

H 2 SO 4 2H + + SO 4 2– ,

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

H 2 SO 4 (diluted) + Cu reaction does not occur,

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

CaCO 3 + H 2 SO 4 = CaSO 4 + H 2 O + CO 2.

Concentrated sulfuric acid is a strong oxidizing agent, especially when heated. It oxidizes many metals, non-metals, and some organic substances. Iron, gold and platinum group metals do not oxidize under the influence of concentrated sulfuric acid (however, iron dissolves well when heated in moderately concentrated sulfuric acid with a mass fraction of 70%). When concentrated sulfuric acid reacts with other metals, sulfates and sulfuric acid reduction products are formed.

2H 2 SO 4 (conc.) + Cu = CuSO 4 + SO 2 + 2H 2 O,

5H 2 SO 4 (conc.) + 8Na = 4Na 2 SO 4 + H 2 S + 4H 2 O,

H 2 SO 4 (conc.) passivates Fe, Al.

When interacting with non-metals, concentrated sulfuric acid is reduced to SO 2:

5H 2 SO 4 (conc.) + 2P = 2H 3 PO 4 + 5SO 2 + 2H 2 O,

2H 2 SO 4 (conc.) + C = 2H 2 O + CO 2 + 2SO 2.

Contact method of receipt sulfuric acid consists of three stages:

1) pyrite firing:

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

2) oxidation of SO 2 to SO 3 in the presence of a catalyst – vanadium oxide:

3) dissolving SO 3 in sulfuric acid to obtain oleum:

SO 3 + H 2 O = H 2 SO 4 + Q,

n SO 3 + H 2 SO 4 (conc.) = H 2 SO 4 n SO3.

Qualitative reaction to sulfate ion– interaction with the barium cation, resulting in the precipitation of a white precipitate, BaSO 4 .

Ba 2+ + SO 4 2– -> BaSO 4,

BaCl 2 + Na 2 SO 4 = BaSO 4 + 2NaCl.

Test on the topic “Sulfur and its compounds”

1. Sulfur and oxygen are:

a) good conductors of electricity;

b) belong to the subgroup of chalcogens;

c) highly soluble in water;

d) have allotropic modifications.

2. As a result of the reaction of sulfuric acid with copper, you can get:

a) hydrogen; b) sulfur;

c) sulfur dioxide; d) hydrogen sulfide.

3. Hydrogen sulfide is:

a) poisonous gas;

b) strong oxidizing agent;

c) typical reducing agent;

d) one of the allotropes of sulfur.

4. The mass fraction (in %) of oxygen in sulfuric anhydride is equal to:

a) 50; b) 60; c) 40; d) 94.

5. Sulfur(IV) oxide is an anhydride:

a) sulfuric acid;

b) sulfurous acid;

c) hydrogen sulfide acid;

d) thiosulfuric acid.

6. By what percentage will the mass of potassium hydrosulfite decrease after calcination?

c) potassium hydrosulfite is thermally stable;

7. You can shift the equilibrium towards the direct reaction of oxidation of sulfur dioxide into sulfuric anhydride:

a) using a catalyst;

b) increasing pressure;

c) reducing pressure;

d) reducing the concentration of sulfur oxide (VI).

8. When preparing a solution of sulfuric acid, you must:

a) pour acid into water;

b) pour water into the acid;

c) the order of infusion does not matter;

d) sulfuric acid does not dissolve in water.

9. What mass (in g) of sodium sulfate decahydrate must be added to 100 ml of 8% sodium sulfate solution (density 1.07 g/ml) to double the mass fraction of salt in the solution?

a) 100; b) 1.07; c) 30.5; d) 22.4.

10. To determine the sulfite ion in qualitative analysis, you can use:

a) lead cations;

b) “iodine water”;

c) solution of potassium permanganate;

d) strong mineral acids.

Key to the test


b, d	V	a, c	b	b	G	b, d	A	V	b, d

Tasks and exercises on sulfur and its compounds

Chain of transformation

1. Sulfur -> iron(II) sulfide -> hydrogen sulfide -> sulfur dioxide -> sulfur trioxide > sulfuric acid > sulfur(IV) oxide.

3. Sulfuric acid -> sulfur dioxide -> sulfur -> sulfur dioxide -> sulfur trioxide -> sulfuric acid.

4. Sulfur dioxide -> sodium sulfite -> sodium hydrosulfite -> sodium sulfite -> sodium sulfate.

5. Pyrite -> sulfur dioxide -> sulfur dioxide -> sulfuric acid -> sulfur oxide (IV) -> potassium sulfite -> sulfur dioxide.

6. Pyrite > sulfur dioxide -> sodium sulfite -> sodium sulfate -> barium sulfate -> barium sulfide.

7. Sodium sulfide -> A -> B -> C -> D -> barium sulfate (all substances contain sulfur; the first, second and fourth reactions are ORR).

Level A

1. 6.5 liters of hydrogen sulfide were passed through a solution containing 5 g of sodium hydroxide. Determine the composition of the resulting solution.

Answer. 7 g NaHS, 5.61 g H2S.

2. What mass of Glauber's salt must be added to 100 ml of 8% sodium sulfate solution (the density of the solution is 1.07 g/ml) to double the mass fraction of the substance in the solution?

Answer. 30.5 g Na 2 SO 4 10H 2 O.

3. To 40 g of a 12% sulfuric acid solution, 4 g of sulfuric anhydride was added. Calculate the mass fraction of the substance in the resulting solution.

Answer. 22% H2SO4.

4. A mixture of iron(II) sulfide and pyrite, weighing 20.8 g, was subjected to prolonged firing, resulting in the formation of 6.72 liters of gaseous product (o.s.). Determine the mass of the solid residue formed during firing.

Answer. 16 g Fe 2 O 3.

5. There is a mixture of copper, carbon and iron(III) oxide with a molar ratio of components of 4:2:1 (in the order listed). What volume of 96% sulfuric acid (density 1.84 g/ml) is needed to completely dissolve 2.2 g of such a mixture when heated?

Answer. 4.16 ml of H 2 SO 4 solution.

6. To oxidize 3.12 g of alkali metal hydrosulfite, it was necessary to add 50 ml of a solution in which the molar concentrations of sodium dichromate and sulfuric acid are 0.2 mol/l and 0.5 mol/l, respectively. Determine the composition and mass of the residue that will be obtained when the solution is evaporated after the reaction.

Answer. 7.47 g mixture of chromium sulfates (3.92 g) and sodium (3.55 g).

Level B

(problems on oleum)

1. What mass of sulfur trioxide must be dissolved in 100 g of 91% sulfuric acid solution to obtain 30% oleum?

Solution

According to the problem:

m(H 2 SO 4) = 100 0.91 = 91 g,

m(H 2 O) = 100 0.09 = 9 g,

(H 2 O) = 9/18 = 0.5 mol.

Part of added SO3 ( m 1) will react with H 2 O:

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

According to the reaction equation:

(SO 3) = (H 2 O) = 0.5 mol.

m 1 (SO 3) = 0.5 80 = 40 g.

Second part SO 3 ( m 2) will be used to create a concentration of oleum. Let us express the mass fraction of oleum:

m 2 (SO 3) = 60 g.

Total mass of sulfur trioxide:

m(SO 3) = m 1 (SO 3) + m 2 (SO 3) = 40 + 60 = 100 g.

Answer. 100 g SO 3.

2. What mass of pyrite must be taken to obtain such an amount of sulfur(VI) oxide that, dissolving it in 54.95 ml of a 91% sulfuric acid solution (density equal to 1.82 g/cm 3), obtain 12.5% oleum? The yield of sulfuric anhydride is considered to be 75%.

Answer. 60 g FeS 2.

3. To neutralize 34.5 g of oleum, 74.5 ml of a 40% solution of potassium hydroxide (density 1.41 g/ml) is consumed. How many moles of sulfuric anhydride are there per 1 mole of sulfuric acid in this oleum?

Answer. 0.5 mol SO3.

4. By adding sulfur(VI) oxide to 300 g of 82% sulfuric acid solution, oleum with a mass fraction of sulfur trioxide of 10% is obtained. Find the mass of sulfuric anhydride used.

Answer. 300 g SO 3.

5. By adding 400 g of sulfur trioxide to 720 g of an aqueous solution of sulfuric acid, oleum with a mass fraction of 7.14% was obtained. Find the mass fraction of sulfuric acid in the original solution.

Answer. 90% H2SO4.

6. Find the mass of a 64% sulfuric acid solution if adding 100 g of sulfur trioxide to this solution produces oleum containing 20% sulfur trioxide.

Answer. 44.4 g of H 2 SO 4 solution.

7. What masses of sulfur trioxide and 91% sulfuric acid solution must be mixed to obtain 1 kg of 20% oleum?

Answer. 428.6 g SO 3 and 571.4 g H 2 SO 4 solution.

8. To 400 g of oleum containing 20% sulfur trioxide, 100 g of a 91% sulfuric acid solution was added. Find the mass fraction of sulfuric acid in the resulting solution.

Answer. 92% H 2 SO 4 in oleum.

9. Find the mass fraction of sulfuric acid in the solution obtained by mixing 200 g of 20% oleum and 200 g of 10% sulfuric acid solution.

Answer. 57.25% H2SO4.

10. What mass of 50% sulfuric acid solution must be added to 400 g of 10% oleum to obtain an 80% sulfuric acid solution?

Answer. 296.67 g of 50% H 2 SO 4 solution.

Answer. 114.83 g oleum.

QUALITATIVE TASKS

1. Colorless gas A with a strong characteristic odor is oxidized by oxygen in the presence of a catalyst into compound B, which is a volatile liquid. Substance B, combining with quicklime, forms salt C. Identify the substances, write the reaction equations.

Answer. Substances: A – SO 2, B – SO 3, C – CaSO 4.

2. When a solution of salt A is heated, precipitate B is formed. The same precipitate is formed when an alkali acts on a solution of salt A. When an acid acts on salt A, gas C is released, which discolors the solution of potassium permanganate. Identify substances, write reaction equations.

Answer. Substances: A – Ca(HSO 3) 2, B – CaSO 3, C – SO 2.

3. When gas A is oxidized with concentrated sulfuric acid, a simple substance B, a complex substance C and water are formed. Solutions of substances A and C react with each other to form a precipitate of substance B. Identify the substances, write the reaction equations.

Answer. Substances: A – H 2 S, B – S, C – SO 2.

4. In the reaction of combining two oxides A and B, liquid at ordinary temperatures, substance C is formed, a concentrated solution of which chars sucrose. Identify substances, write reaction equations.

Answer. Substances: A – SO 3, B – H 2 O, C – H 2 SO 4.

5. At your disposal are iron(II) sulfide, aluminum sulfide and aqueous solutions of barium hydroxide and hydrogen chloride. Obtain seven different salts from these substances (without using ORR).

Answer. Salts: AlCl 3, BaS, FeCl 2, BaCl 2, Ba(OH)Cl, Al(OH)Cl 2, Al(OH) 2 Cl.

6. When concentrated sulfuric acid acts on bromides, sulfur dioxide is released, and on iodides, hydrogen sulfide is released. Write the reaction equations. Explain the difference in the nature of the products in these cases.

Answer. Reaction equations:

2H 2 SO 4 (conc.) + 2NaBr = SO 2 + Br 2 + Na 2 SO 4 + 2H 2 O,

5H 2 SO 4 (conc.) + 8NaI = H 2 S + 4I 2 + 4Na 2 SO 4 + 4H 2 O.

1 See: Lidin R.A."Handbook of general and inorganic chemistry". M.: Education, 1997.

* The +/– sign means that this reaction does not occur with all reagents or under specific conditions.

To be continued

O.S.ZAYTSEV

CHEMISTRY BOOK

FOR SECONDARY SCHOOL TEACHERS,
STUDENTS OF PEDAGOGICAL UNIVERSITIES AND SCHOOLCHILDREN OF 9–10 GRADES,
WHO HAVE 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 ultraviolet rays of the Sun, 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.

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 ozonator. 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.

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

3. Reductive properties of hydrogen sulfide and sulfide ion.

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 the 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.

3b. Reaction of hydrogen sulfide with sulfuric acid.

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 produced and worked with under draft.

4a. Interaction of sulfurous acid with hydrogen peroxide.

4b. Reaction between sulfur dioxide and hydrogen sulfide.

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

Obtaining hydrogen sulfide. Production of sulfur dioxide by burning sulfur, hydrogen sulfide and other types of raw materials Hydrogen sulfide sulfur dioxide

§ 8.1. Redox reactions

Chemistry tutor

O.S.ZAYTSEV

CHEMISTRY BOOK

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

§ 8.1. Redox reactions

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