Sulfurous acid is an inorganic dibasic unstable acid of medium strength. An unstable compound, known only in aqueous solutions at a concentration of no more than six percent. When attempting to isolate pure sulfurous acid, it breaks down into sulfur oxide (SO2) and water (H2O). For example, when concentrated sulfuric acid (H2SO4) reacts with sodium sulfite (Na2SO3), sulfur oxide (SO2) is released instead of sulfurous acid. This is what the reaction looks like:
Na2SO3 (sodium sulfite) + H2SO4 ( sulfuric acid) = Na2SO4 (sodium sulfate) + SO2 (sulfur dioxide) + H2O (water)
Sulfurous acid solution
When storing it, it is necessary to exclude access to air. Otherwise, sulfurous acid, slowly absorbing oxygen (O2), will turn into sulfuric acid.
2H2SO3 (sulfuric acid) + O2 (oxygen) = 2H2SO4 (sulfuric acid)
Solutions of sulfurous acid have a rather specific odor (reminiscent of the odor remaining after lighting a match), the presence of which can be explained by the presence of sulfur oxide (SO2), which is not chemically bound with water.
Chemical properties sulfurous acid
1. H2SO3) can be used as a reducing agent or an oxidizing agent.
H2SO3 is a good reducing agent. With its help, it is possible to obtain hydrogen halides from free halogens. For example:
H2SO3 (sulfuric acid) + Cl2 (chlorine, gas) + H2O (water) = H2SO4 (sulfuric acid) + 2HCl ( hydrochloric acid)
But when interacting with strong reducing agents, this acid will act as an oxidizing agent. An example is the reaction of sulfurous acid with hydrogen sulfide:
H2SO3 (sulfuric acid) + 2H2S (hydrogen sulfide) = 3S (sulfur) + 3H2O (water)
2. The chemical compound we are considering forms two - sulfites (medium) and hydrosulfites (acidic). These salts are reducing agents, just like (H2SO3) sulfurous acid. When they are oxidized, salts of sulfuric acid are formed. When sulfites of active metals are calcined, sulfates and sulfides are formed. This is a self-oxidation-self-healing reaction. For example:
4Na2SO3 (sodium sulfite) = Na2S + 3Na2SO4 (sodium sulfate)
Sodium and potassium sulfites (Na2SO3 and K2SO3) are used in dyeing fabrics in the textile industry, in bleaching metals, and in photography. Calcium hydrosulfite (Ca(HSO3)2), which exists only in solution, is used to process wood material into a special sulfite pulp. It is then used to make paper.
Application of sulfurous acid
Sulfurous acid is used:
For bleaching wool, silk, wood pulp, paper and other similar substances that cannot withstand bleaching with stronger oxidizing agents (for example, chlorine);
As a preservative and antiseptic, for example, to prevent the fermentation of grain when producing starch or to prevent the fermentation process in wine barrels;
To preserve food, for example, when canning vegetables and fruits;
Processed into sulfite pulp, from which paper is then produced. In this case, a solution of calcium hydrosulfite (Ca(HSO3)2) is used, which dissolves lignin, a special substance that binds cellulose fibers.
Sulfurous acid: preparation
This acid can be produced by dissolving sulfur dioxide (SO2) in water (H2O). You will need concentrated sulfuric acid (H2SO4), copper (Cu) and a test tube. Algorithm of actions:
1. Carefully pour concentrated sulfuric acid into a test tube and then place a piece of copper in it. Heat up. The following reaction occurs:
Cu (copper) + 2H2SO4 (sulfuric acid) = CuSO4 (sulfur sulfate) + SO2 (sulfur dioxide) + H2O (water)
2. The flow of sulfur dioxide must be directed into a test tube with water. When it dissolves, it partially occurs with water, resulting in the formation of sulfurous acid:
SO2 (sulfur dioxide) + H2O (water) = H2SO3
So, by passing sulfur dioxide through water, you can get sulfurous acid. It is worth considering that this gas has an irritating effect on the membranes of the respiratory tract, can cause inflammation, as well as loss of appetite. Inhaling it for a long time may cause loss of consciousness. This gas must be handled with extreme caution and care.
Hydrogen sulfide (H₂S) is a colorless gas with a rotten egg odor. It is denser than hydrogen. Hydrogen sulfide is deadly poisonous to humans and animals. Even a small amount of it in the air causes dizziness and nausea, but the worst thing is that after inhaling it for a long time, this smell is no longer felt. However, for hydrogen sulfide poisoning, there is a simple antidote: you should wrap a piece of bleach in a handkerchief, then moisten it, and sniff the package for a while. Hydrogen sulfide is produced by reacting sulfur with hydrogen at a temperature of 350 °C:
H₂ + S → H₂S
This is a redox reaction: during it, the oxidation states of the elements participating in it change.
In laboratory conditions, hydrogen sulfide is produced by treating iron sulfide with sulfuric or hydrochloric acid:
FeS + 2HCl → FeCl₂ + H₂S
This is an exchange reaction: in it, the interacting substances exchange their ions. This process is usually performed using a Kipp apparatus.
Kipp apparatus
Properties of hydrogen sulfide
When hydrogen sulfide burns, sulfur oxide 4 and water vapor are formed:
2H₂S + 3О₂ → 2Н₂О + 2SO₂
H₂S burns with a bluish flame, and if you hold an inverted beaker over it, clear condensate (water) will appear on its walls.
However, with a slight decrease in temperature, this reaction proceeds somewhat differently: a yellowish coating of free sulfur will appear on the walls of the pre-cooled glass:
2H₂S + O₂ → 2H₂O + 2S
The industrial method for producing sulfur is based on this reaction.
When a pre-prepared gaseous mixture of hydrogen sulfide and oxygen is ignited, an explosion occurs.
The reaction of hydrogen sulfide and sulfur(IV) oxide also produces free sulfur:
2H₂S + SO₂ → 2H₂O + 3S
Hydrogen sulfide is soluble in water, and three volumes of this gas can dissolve in one volume of water, forming weak and unstable hydrosulfide acid (H₂S). This acid is also called hydrogen sulfide water. As you can see, the formulas of hydrogen sulfide gas and hydrogen sulfide acid are written the same way.
If a solution of lead salt is added to hydrosulfide acid, a black precipitate of lead sulfide will form:
H₂S + Pb(NO₃)₂ → PbS + 2HNO₃
This is a qualitative reaction for the detection of hydrogen sulfide. It also demonstrates the ability of hydrosulfide acid to enter into exchange reactions with salt solutions. Thus, any soluble lead salt is a reagent for hydrogen sulfide. Some other metal sulfides also have a characteristic color, for example: zinc sulfide ZnS - white, cadmium sulfide CdS - yellow, copper sulfide CuS - black, antimony sulfide Sb₂S₃ - red.
By the way, hydrogen sulfide is an unstable gas and, when heated, almost completely decomposes into hydrogen and free sulfur:
H₂S → H₂ + S
Hydrogen sulfide interacts intensively with aqueous solutions of halogens:
H₂S + 4Cl₂ + 4H₂O→ H₂SO₄ + 8HCl
Hydrogen sulfide in nature and human activity
Hydrogen sulfide is part of volcanic gases, natural gas and gases associated with oil fields. There is also a lot of it in natural mineral waters, for example, in the Black Sea it lies at a depth of 150 meters and below.
Hydrogen sulfide is used:
- in medicine (treatment with hydrogen sulfide baths and mineral waters);
- in industry (production of sulfur, sulfuric acid and sulfides);
- in analytical chemistry (for the precipitation of heavy metal sulfides, which are usually insoluble);
- in organic synthesis (to produce sulfur analogues of organic alcohols (mercaptans) and thiophene (sulphur-containing aromatic hydrocarbon). Another recently emerging area in science is hydrogen sulfide energy. The production of energy from hydrogen sulfide deposits from the bottom of the Black Sea is being seriously studied.
The nature of redox reactions of sulfur and hydrogen
The reaction of hydrogen sulfide formation is redox:
Н₂⁰ + S⁰→ H₂⁺S²⁻
The process of interaction of sulfur with hydrogen is easily explained by the structure of their atoms. Hydrogen occupies first place in the periodic table, therefore, the charge of its atomic nucleus is equal to (+1), and 1 electron circles around the atomic nucleus. Hydrogen easily gives up its electron to atoms of other elements, turning into a positively charged hydrogen ion - a proton:
Н⁰ -1е⁻= Н⁺
Sulfur is in position sixteen in the periodic table. This means that the charge of the nucleus of its atom is (+16), and the number of electrons in each atom is also 16e⁻. The location of sulfur in the third period suggests that its sixteen electrons swirl around the atomic nucleus, forming 3 layers, the last of which contains 6 valence electrons. The number of valence electrons of sulfur corresponds to the number of group VI in which it is located in the periodic table.
So, sulfur can donate all six valence electrons, as in the case of the formation of sulfur(VI) oxide:
2S⁰ + 3O2⁰ → 2S⁺⁶O₃⁻²
In addition, as a result of the oxidation of sulfur, 4e⁻ can be given up by its atom to another element to form sulfur(IV) oxide:
S⁰ + O2⁰ → S⁺4 O2⁻²
Sulfur can also donate two electrons to form sulfur(II) chloride:
S⁰ + Cl2⁰ → S⁺² Cl2⁻
In all three of the above reactions, sulfur donates electrons. Consequently, it is oxidized, but at the same time acts as a reducing agent for oxygen atoms O and chlorine Cl. However, in the case of the formation of H2S, oxidation is the lot of hydrogen atoms, since they are the ones who lose electrons, restoring the external energy level sulfur from six electrons to eight. As a result, each hydrogen atom in its molecule becomes a proton:
Н2⁰-2е⁻ → 2Н⁺,
and the sulfur molecule, on the contrary, being reduced, turns into a negatively charged anion (S⁻²): S⁰ + 2е⁻ → S⁻²
Thus, in chemical reaction In the formation of hydrogen sulfide, the oxidizing agent is sulfur.
From the point of view of the manifestation of sulfur in various oxidation states, another interesting interaction between sulfur(IV) oxide and hydrogen sulfide is the reaction to produce free sulfur:
2H₂⁺S-²+ S⁺⁴О₂-²→ 2H₂⁺O-²+ 3S⁰
As can be seen from the reaction equation, both the oxidizing agent and the reducing agent in it are sulfur ions. Two sulfur anions (2-) donate two of their electrons to the sulfur atom in the sulfur(II) oxide molecule, as a result of which all three sulfur atoms are reduced to free sulfur.
2S-² - 4е⁻→ 2S⁰ - reducing agent, oxidizes;
S⁺⁴ + 4е⁻→ S⁰ - oxidizing agent, reduced.
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:
and how 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 section “H 2 S”). 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 during roasting of 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).
Used in the production of sulfates and other sulfur compounds, mineral fertilizers, explosives, dyes and medicines, in organic synthesis, for the “opening” (the first stage of processing) of industrially important ores and minerals, during the purification of petroleum products, the electrolysis of water, as an electrolyte for lead batteries. Poisonous, 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)
(SO 3 absorption clean water with the production of H 2 SO 4 is not carried out due to the strong heating of the mixture and 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 (dark purple hydrate, 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 iron sulfate), 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 Nv 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:
Reversible reaction Ammonia production 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)
Used for ammonia synthesis, 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 and 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... HOH. 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).
Related information.
Oxygen with the most common element in the earth's crust. The oxygen molecule is diatomic (O 2). A simple substance - molecular oxygen - is a colorless and odorless gas, poorly soluble in water. The Earth's atmosphere contains 21% (by volume) oxygen. In natural compounds, oxygen occurs in the form of oxides (H 2 O, SiO 2) and salts of oxoacids. One of the most important natural oxygen compounds is water, or hydrogen oxide H2O.
In addition to oxides, oxygen is capable of forming peroxides - substances containing the following group of atoms: –O–O–. One of the most important peroxides is hydrogen peroxide H 2 O 2 (H–O–O–H). In peroxides, oxygen atoms have an intermediate oxidation state of minus 1, so these compounds can be both oxidizing and reducing agents:
From the values of standard electrode potentials it follows that oxides
The thermal properties of H2O2 are most pronounced in an acidic environment, and the reducing properties are most pronounced in an alkaline environment. For example, hydrogen peroxide in an acidic environment is capable of oxidizing those substances whose standard potential of the electrochemical system does not exceed +1.776 V, and reducing only those whose potential is greater than +0.682 V.
An allotropic modification of oxygen is ozone (O3), a gas with a specific odor. Ozone is produced by the action of “quiet” electrical discharges on oxygen in special devices - ozonizers. The reaction of converting oxygen into ozone requires energy:
3O2 ↔ 2O3 – 285 kJ.
The reverse process—ozone decomposition—occurs spontaneously.
Ozone is one of the strongest oxidizing agents; in terms of oxidative activity it is second only to fluorine.
At high temperatures, sulfur reacts with hydrogen to form hydrogen sulfide(H2S) is a colorless gas with a characteristic odor of rotting protein. Since this reaction is reversible, in practice hydrogen sulfide is usually produced by the action of dilute acids on metal sulfides:
FeS + 2 HCl → H2S + FeCl2.
Hydrogen sulfide is a strong reducing agent; When ignited in air, it burns with a bluish flame:
2 H2S + 3 O2 → 2 SO2 + 2 H2O (in excess oxygen).
Therefore, a mixture of hydrogen sulfide with air is explosive. With a lack of oxygen, hydrogen sulfide is oxidized only to free sulfur:
2 H2S + O2 → 2 S + 2 H2O.
Hydrogen sulfide is very poisonous and can cause severe poisoning.
A solution of hydrogen sulfide in water has the properties of a weak dibasic acid (K1 = 6×10–8, K2 = 1×10–14). Medium salts of hydrosulfide acid - sulfides - can be obtained by direct interaction of metals with sulfur. Slightly soluble sulfides can be obtained by reacting hydrogen sulfide with solutions of salts of the corresponding metals:
CuSO4 + H2S CuS+ H2SO4 .
Sulfur oxide(IV) is formed when sulfur burns in air:
S + O2 → SO2.
In industry, SO2 is obtained by roasting metal sulfides and polysulfides, as well as by thermal decomposition of sulfates (in particular CaSO4):
Sulfur dioxide is a colorless gas with the smell of burnt sulfur. SO2 dissolves well in water, forming sulfurous acid:
Sulfurous acid– weak dibasic acid (K1=1.6×10–2, K2=6×10–8). H2SO3 and its salts are good reducing agents and are oxidized to sulfuric acid or sulfates:
At high temperatures in the presence of a catalyst (V2O5, platinum-based alloys), sulfur dioxide is oxidized by oxygen to trioxide:
Sulfur (VI) oxide is sulfuric acid anhydride:
In the gaseous state, sulfur oxide (VI) consists of SO3 molecules arranged in the shape of a regular triangle. When SO3 vapor condenses, a volatile liquid is formed (boiling point = +44.8 °C), consisting mainly of trimeric cyclic molecules. When cooled to +16.8 °C, it solidifies and the so-called ice-like modification SO3 is formed. During storage, it gradually turns into an asbestos-like modification of SO3, consisting of polymer molecules.
Concentrated sulfuric acid, especially hot, is a vigorous oxidizing agent. It oxidizes bromide and iodide ions to free halogens, coal to carbon dioxide, and sulfur to SO2. When interacting with metals, concentrated sulfuric acid converts them into sulfates, reducing them to SO2, S or H2S. The more active the metal, the more deeply the acid is reduced.
For example, when concentrated sulfuric acid reacts with copper, SO2 is predominantly released; When interacting with zinc, simultaneous release of sulfur (IV) oxide, free sulfur, and hydrogen sulfide can be observed:
H2SO4 is a strong dibasic acid, dissociated in the first stage
almost completely; dissociation in the second stage occurs to a lesser extent, however, in dilute aqueous solutions, sulfuric acid is dissociated almost completely according to the following scheme:
H2SO4 → 2 H + + SO4 2-
Most sulfuric acid salts are highly soluble in water. The practically insoluble ones include BaSO4, SrSO4, PbSO4; slightly soluble CaSO4. The qualitative reaction to SO4 2– ions is due to the formation of poorly soluble sulfates. For example, when barium ions are introduced into a solution containing sulfathions, a white precipitate of barium sulfate precipitates, practically insoluble in water and dilute acids:
Ba 2+ + SO4 2- → BaSO4↓ .
Sulfuric acid is used in the production of mineral fertilizers;
as an electrolyte in lead batteries; for obtaining various mineral acids and salts; in the production of chemical fibers, dyes, smoke-forming and explosives; in the oil, metalworking, textile, leather and other industries, etc.
Almurzinova Zavrish Bisembaevna , teacher of biology and chemistry MBOU "State Farm Basic secondary school 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. Sulfur 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.
Lesson progress :
Organizational moment .
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) at long-term storage 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, based on? river sand?
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? We find out in conversation physical properties, we note the physiological effect. We determine the chemical properties by experimenting with the combustion of hydrogen sulfide in air at different 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)