Hydrogen sulfide with sulfur dioxide. Chemistry tutor manual. Coking of hard coals
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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 (valence 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. Has four stable isotopes with mass numbers 32, 33, 34 and 36. Natural sulfur is 95% composed of the isotope 32 S.
Russian name 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, then plastic sulfur is formed, its properties reminiscent of rubber. Plastic sulfur consists of long zigzag chains. This allotropic modification is unstable and spontaneously transforms into one of the crystalline forms.
Orthorhombic sulfur is a crystalline solid yellow color; 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
According to their own chemical properties sulfur is a typical active nonmetal. 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 blood hemoglobin, 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.
Portion 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
Sulfur oxide (sulfur dioxide, sulfur dioxide, sulfur dioxide) is a colorless gas that under normal conditions has a sharp characteristic odor (similar to the smell of a burning match). Liquefied under pressure room temperature. Sulfur dioxide is soluble in water, and unstable sulfuric acid is formed. This substance is also soluble in sulfuric acid and ethanol. This is one of the main components that make up volcanic gases.
1. Sulfur dioxide dissolves in water, resulting in sulfurous acid. Under normal conditions, this reaction is reversible.
SO2 (sulfur dioxide) + H2O (water) = H2SO3 (sulfurous acid).
2. With alkalis, sulfur dioxide forms sulfites. For example: 2NaOH (sodium hydroxide) + SO2 (sulfur dioxide) = Na2SO3 (sodium sulfite) + H2O (water).
3. The chemical activity of sulfur dioxide is quite high. The reducing properties of sulfur dioxide are most pronounced. In such reactions, the oxidation state of sulfur increases. For example: 1) SO2 (sulfur dioxide) + Br2 (bromine) + 2H2O (water) = H2SO4 (sulfuric acid) + 2HBr (hydrogen bromide); 2) 2SO2 (sulfur dioxide) + O2 (oxygen) = 2SO3 (sulfite); 3) 5SO2 (sulfur dioxide) + 2KMnO4 (potassium permanganate) + 2H2O (water) = 2H2SO4 (sulfuric acid) + 2MnSO4 (manganese sulfate) + K2SO4 (potassium sulfate).
The last reaction is an example of a qualitative reaction to SO2 and SO3. The solution becomes purple in color.)
4. In the presence of strong reducing agents, sulfur dioxide can exhibit oxidizing properties. For example, in order to extract sulfur from exhaust gases in the metallurgical industry, they use the reduction of sulfur dioxide with carbon monoxide (CO): SO2 (sulfur dioxide) + 2CO (carbon monoxide) = 2CO2 + S (sulfur).
Also, the oxidizing properties of this substance are used to obtain phosphorous acid: PH3 (phosphine) + SO2 (sulfur dioxide) = H3PO2 (phosphoric acid) + S (sulfur).
Where is sulfur dioxide used?
Sulfur dioxide is mainly used to produce sulfuric acid. It is also used in the production of low-alcohol drinks (wine and other mid-price drinks). Due to the property of this gas to kill various microorganisms, it is used to fumigate warehouses and vegetable stores. In addition, sulfur oxide is used to bleach wool, silk, and straw (those materials that cannot be bleached with chlorine). In laboratories, sulfur dioxide is used as a solvent and in order to obtain various salts of sulfur dioxide.
Physiological effects
Sulfur dioxide has strong toxic properties. Symptoms of poisoning are cough, runny nose, hoarseness, a peculiar taste in the mouth, and severe sore throat. When sulfur dioxide is inhaled in high concentrations, difficulty swallowing and choking, speech disturbance, nausea and vomiting occur, and acute pulmonary edema may develop.
MPC of sulfur dioxide:
- indoors - 10 mg/m³;
- average daily maximum one-time exposure in atmospheric air - 0.05 mg/m³.
Sensitivity to sulfur dioxide varies among individuals, plants, and animals. For example, among trees the most resistant are oak and birch, and the least resistant are spruce and pine.
Sulfur– element of the 3rd period and VIA group Periodic table, serial number 16, refers to chalcogens. The electronic formula of the atom is [ 10 Ne]3s 2 3p 4, the characteristic oxidation states are 0, -II, +IV and +VI, the S VI state is considered stable.
Scale of sulfur oxidation states:
The electronegativity of sulfur is 2.60 and is characterized by non-metallic properties. In hydrogen and oxygen compounds it is found in various anions and forms oxygen-containing acids and their salts, binary compounds.
In nature - fifteenth element by chemical abundance (seventh among non-metals). It is found in free (native) and bound form. A vital element for higher organisms.
Sulfur S. Simple substance. Yellow crystalline (α‑rhombic and β‑monoclinic,
at 95.5 °C) or amorphous (plastic). At the nodes of the crystal lattice there are S 8 molecules (non-planar rings of the “crown” type), amorphous sulfur consists of S n chains. A low-melting substance, the viscosity of the liquid passes through a maximum at 200 °C (breakdown of S 8 molecules, interweaving of S n chains). The pair contains molecules S 8, S 6, S 4, S 2. At 1500 °C, monoatomic sulfur appears (in chemical equations, for simplicity, any sulfur is depicted as S).
Sulfur is insoluble in water and under normal conditions does not react with it; it is highly soluble in carbon disulfide CS 2.
Sulfur, especially powdered sulfur, is highly active when heated. Reacts as an oxidizing agent with metals and non-metals:
but as reducing agent– with fluorine, oxygen and acids (boiling):
Sulfur undergoes dismutation in alkali solutions:
3S 0 + 6KOH (conc.) = 2K 2 S ‑II + K 2 S IV O 3 + 3H 2 O
At high temperatures (400 °C), sulfur displaces iodine from hydrogen iodide:
S + 2HI (g) = I 2 + H 2 S,
but in solution the reaction goes in the opposite direction:
I 2 + H 2 S (p) = 2 HI + S↓
Receipt: V industry smelted from natural deposits of native sulfur (using water vapor), released during desulfurization of coal gasification products.
Sulfur is used for the synthesis of carbon disulfide, sulfuric acid, sulfur (vat) dyes, during the vulcanization of rubber, as a means of protecting plants from powdery mildew, 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 when roasting pyrite):
S + O 2 = SO 2(280–360 °C)
4FeS 2 + 11O 2 = 2Fe 2 O 3 + 8 SO 2(800 °C, firing)
in the laboratory - displacement of sulfites with sulfuric acid:
BaSO 3 (t) + H 2 SO 4 (conc.) = BaSO 4 ↓ + SO 2 + H 2 O
Sodium sulfite Na 2 SO 3. Oxosol. White. When heated in air, it decomposes without melting and melts under excess pressure of argon. When wet and in solution, it is sensitive to atmospheric oxygen. It is highly soluble in water and hydrolyzes at the anion. Decomposes with acids. Typical reducer.
Qualitative reaction on the SO 3 2‑ ion - the formation of a white precipitate of barium sulfite, which is transferred into solution with strong acids (HCl, HNO 3).
It is used as a reagent in analytical chemistry, a component of photographic solutions, and a chlorine neutralizer for bleaching fabrics.
Equations of the most important reactions:
Receipt:
Na 2 CO 3 (conc.) + SO 2 = Na2SO3+CO2
Sulfuric acid. Sulfates
Sulfuric acid H 2 SO 4. Oxoacid. Colorless liquid, very viscous (oily), very hygroscopic. The molecule has a distorted tetrahedral structure (sp 3 -hybridization), contains covalent σ-bonds S – OH and σπ-bonds S=O. The SO 4 2‑ ion has a regular tetrahedral structure. It has a wide temperature range of the liquid state (~300 degrees). Partially decomposes when heated above 296 °C. It is distilled in the form of an azeotropic mixture with water (mass fraction of acid 98.3%, boiling point 296–340 °C), and with stronger heating it decomposes completely. Unlimitedly miscible with water (with strong exo‑effect). Strong acid in solution, neutralized by alkalis and ammonia hydrate. Converts metals into sulfates (with an excess of concentrated acid under normal conditions, soluble hydrosulfates are formed), but the metals Be, Bi, Co, Fe, Mg and Nb are passivated in concentrated acid and do not react with it. Reacts with basic oxides and hydroxides, decomposes salts of weak acids. A weak oxidizing agent in a dilute solution (due to H I), a strong oxidizing agent in a concentrated solution (due to S VI). It dissolves SO 3 well and reacts with it (a heavy oily liquid is formed - oleum, contains H 2 S 2 O 7).
Qualitative reaction on the SO 4 2‑ ion – precipitation of white barium sulfate BaSO 4 (the precipitate is not transferred into solution by hydrochloric and nitric acids, unlike the white precipitate BaSO 3).
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. Toxic, causes skin burns. Equations of the most important reactions:
Receipt V industry:
a) synthesis of SO 2 from sulfur, sulfide ores, hydrogen sulfide and sulfate ores:
S + O 2 (air) = SO 2(280–360 °C)
4FeS 2 + 11O 2 (air) = 8 SO 2+ 2Fe 2 O 3 (800 °C, firing)
2H 2 S + 3O 2 (g) = 2 SO 2+ 2H 2 O (250–300 °C)
CaSO 4 + C (coke) = CaO + SO 2+ CO (1300–1500 °C)
b) conversion of SO 2 to SO 3 in a contact apparatus:
c) synthesis of concentrated and anhydrous sulfuric acid:
H 2 O (dil. H 2 SO 4) + SO 3 = H2SO4(conc., anhydrous)
(the absorption of SO 3 with pure water to produce H 2 SO 4 is not carried out due to the strong heating of the mixture and the reverse decomposition of H 2 SO 4, see above);
d) synthesis oleum– a mixture of anhydrous H 2 SO 4, disulfuric acid H 2 S 2 O 7 and excess SO 3. Dissolved SO 3 guarantees the anhydrity of oleum (when water enters, H 2 SO 4 is immediately formed), which allows it to be safely transported in steel tanks.
Sodium sulfate Na 2 SO 4. Oxosol. White, hygroscopic. Melts and boils without decomposition. Forms crystalline hydrate (mineral mirabilite), easily losing water; technical name Glauber's salt. It is highly soluble in water and does not hydrolyze. Reacts with H 2 SO 4 (conc.), SO 3 . It is reduced by hydrogen and coke when heated. Enters into ion exchange reactions.
Used in the production of glass, cellulose and mineral paints, as 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, component 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. Defines most“constant” hardness of fresh water (for more details, see 9.2).
Equations of the most important reactions: 100–128 °C
It is used as a raw material in the production of SO 2, H 2 SO 4 and (NH 4) 2 SO 4, as a flux in metallurgy, and as a paper filler. A binder mortar made from burnt gypsum “sets” faster than a mixture based on Ca(OH) 2 . Hardening is ensured by the binding of water, the formation of gypsum in the form of a stone mass. Burnt gypsum is used to make plaster casts, architectural and decorative forms and products, partition slabs and panels, and stone floors.
Aluminum-potassium sulfate KAl(SO 4) 2. Double oxo salt. White, hygroscopic. Decomposes when heated strongly. Forms crystalline hydrate - potassium alum. Moderately soluble in water, hydrolyzes with aluminum cation. Reacts with alkalis, ammonia hydrate.
It is used as a mordant for dyeing fabrics, a leather tanning agent, a coagulant for purifying fresh water, a component of compositions for sizing paper, and an external hemostatic agent in medicine and cosmetology. It is formed by the joint crystallization of aluminum and potassium sulfates.
Equations of the most important reactions:
Chromium(III) sulfate - potassium KCr(SO 4) 2. Double oxo salt. Red (hydrate dark purple, technical name chromium-potassium alum). When heated, it decomposes without melting. It is highly soluble in water (the gray-blue color of the solution corresponds to aqua complex 3+), hydrolyzes at the chromium(III) cation. Reacts with alkalis, ammonia hydrate. Weak oxidizing and reducing agent. Enters into ion exchange reactions.
Qualitative reactions on the Cr 3+ ion – reduction to Cr 2+ or oxidation to yellow CrO 4 2‑.
It is used as a leather tanning agent, a mordant for dyeing fabrics, and a reagent in photography. It is formed by the joint crystallization of chromium(III) and potassium sulfates. Equations of the most important reactions:
Manganese (II) sulfate MnSO 4 . Oxosol. White, melts and decomposes when heated. Crystalline hydrate MnSO 4 5H 2 O – red-pink, technical name manganese sulfate. It is highly soluble in water; the light pink (almost colorless) color of the solution corresponds to aquacomplex 2+; hydrolyzes at the cation. Reacts with alkalis, ammonia hydrate. Weak reducing agent, reacts with typical (strong) oxidizing agents.
Qualitative reactions on the Mn 2+ ion – commutation with the MnO 4 ion and the disappearance of the violet color of the latter, oxidation of Mn 2+ to MnO 4 and the appearance of a violet color.
It is used for the production of Mn, MnO 2 and other manganese compounds, as a microfertilizer and analytical reagent.
Equations of the most important reactions:
Receipt:
2MnO 2 + 2H 2 SO 4 (conc.) = 2 MnSO4+ O 2 + 2H 2 O (100 °C)
Iron (II) sulfate FeSO 4 . Oxosol. White (light green hydrate, technical name inkstone), hygroscopic. Decomposes when heated. It is highly soluble in water and is slightly hydrolyzed by the cation. It is quickly oxidized in solution by atmospheric oxygen (the solution turns yellow and becomes cloudy). Reacts with oxidizing acids, alkalis, and ammonia hydrate. Typical reducer.
It is used as a component of mineral paints, electrolytes in electroplating, a wood preservative, a fungicide, and a medicine against anemia. In the laboratory it is often taken in the form of a double salt Fe(NH 4) 2 (SO 4) 2 6H 2 O ( Mohr's salt), more resistant to air.
Equations of the most important reactions:
Receipt:
Fe + H 2 SO 4 (diluted) = FeSO4+H2
FeCO 3 + H 2 SO 4 (diluted) = FeSO4+ CO 2 + H 2 O
7.4. Non-metals VA‑group
Nitrogen. Ammonia
Nitrogen– element of the 2nd period and VA group of the Periodic system, serial number 7. Electronic formula of the atom [ 2 He]2s 2 2p 3, characteristic oxidation states 0, ‑III, +III and +V, less often +II, +IV and etc.; the N v state is considered relatively stable.
Scale of nitrogen oxidation states:
Nitrogen has a high electronegativity (3.07), third after F and O. It exhibits typical non-metallic (acidic) properties. Forms various oxygen-containing acids, salts and binary compounds, as well as the ammonium cation NH 4 + and its salts.
In nature - seventeenth by chemical abundance element (ninth among non-metals). A vital element for all organisms.
Nitrogen N 2. Simple substance. It consists of non-polar molecules with a very stable σππ-bond N ≡ N, this explains the chemical inertness of nitrogen under normal conditions. A colorless, tasteless and odorless gas that condenses into a colorless liquid (unlike O2).
Main component of air: 78.09% by volume, 75.52% by mass. Nitrogen boils away from liquid air before oxygen O2. Slightly soluble in water (15.4 ml/1 l H 2 O at 20 ° C), the solubility of nitrogen is less than that of oxygen.
At room temperature, N2 reacts only with lithium (in a humid atmosphere), forming lithium nitride Li3N; nitrides of other elements are synthesized with strong heating:
N 2 + 3Mg = Mg 3 N 2 (800 °C)
In an electrical discharge, N2 reacts with fluorine and, to a very small extent, with oxygen:
The reversible reaction to produce ammonia occurs at 500 °C, under pressure up to 350 atm and always in the presence of a catalyst (Fe/F 2 O 3 /FeO, in the laboratory Pt):
According to Le Chatelier's principle, an increase in ammonia yield should occur with increasing pressure and decreasing temperature. However, the reaction rate at low temperatures is very low, so the process is carried out at 450–500 °C, achieving a 15% ammonia yield. Unreacted N 2 and H 2 are returned to the reactor and thereby increase the degree of reaction.
Nitrogen is chemically passive in relation to acids and alkalis and does not support combustion.
Receipt V industry– fractional distillation of liquid air or removal of oxygen from air by chemical means, for example, by the reaction 2C (coke) + O 2 = 2CO when heated. In these cases, nitrogen is obtained, which also contains impurities of noble gases (mainly argon).
IN laboratories small amounts of chemically pure nitrogen can be obtained by the commutation reaction with moderate heating:
N ‑III H 4 N III O 2(t) = N 2 0 + 2H 2 O (60–70 °C)
NH 4 Cl (p) + KNO 2 (p) = N 2 0 + KCl + 2H 2 O (100 °C)
It is used for the synthesis of ammonia, nitric acid and other nitrogen-containing products, as an inert medium for chemical and metallurgical processes and storage of flammable substances.
Ammonia NH3. Binary compound, the oxidation state of nitrogen is – III. Colorless gas with a sharp characteristic odor. The molecule has the structure of an incomplete tetrahedron [: N(H) 3)] (sp 3 -hybridization). The presence of a donor pair of electrons on the sp 3 -hybrid orbital of nitrogen in the NH 3 molecule determines the characteristic reaction of addition of a hydrogen cation, which results in the formation of a cation ammonium NH4+. It liquefies under excess pressure at room temperature. In the liquid state, it is associated through hydrogen bonds. Thermally unstable. Highly soluble in water (more than 700 l/1 l H 2 O at 20 °C); the proportion in the saturated solution is = 34% by mass and = 99% by volume, pH = 11.8.
Very reactive, prone to addition reactions. Cr reacts in oxygen, reacts with acids. It exhibits reducing (due to N‑III) and oxidizing (due to H I) properties. It is dried only with calcium oxide.
Qualitative reactions– formation of white “smoke” upon contact with gaseous HCl, blackening of a piece of paper moistened with a solution of Hg 2 (NO 3) 2.
An intermediate product in the synthesis of HNO 3 and ammonium salts. Used in the production of soda, nitrogen fertilizers, dyes, explosives; liquid ammonia is a refrigerant. Poisonous.
Equations of the most important reactions:
Receipt: V laboratories– displacement of ammonia from ammonium salts when heated with soda lime (NaOH + CaO):
or boiling an aqueous solution of ammonia and then drying the gas.
IN industry ammonia is synthesized from nitrogen (see) with hydrogen. Produced by industry either in liquefied form or in the form of a concentrated aqueous solution under the technical name ammonia water.
Ammonia hydrate NH 3 H 2 O. Intermolecular connection. White, in the crystal lattice - molecules NH 3 and H 2 O, connected by a weak hydrogen bond H 3 N ... HON. Present in an aqueous solution of ammonia, a weak base (dissociation products - NH 4 ‑ cation and OH ‑ anion). The ammonium cation has a regular tetrahedral structure (sp 3 hybridization). Thermally unstable, completely decomposes when the solution is boiled. Neutralized by strong acids. Shows reducing properties (due to N III) in a concentrated solution. Enters into ion exchange and complexation reactions.
Qualitative reaction– formation of white “smoke” upon contact with gaseous HCl.
It is used to create a slightly alkaline environment in solution during the precipitation of amphoteric hydroxides.
A 1M ammonia solution contains mainly NH 3 H 2 O hydrate and only 0.4% NH 4 + and OH - ions (due to hydrate dissociation); Thus, the ionic “ammonium hydroxide NH 4 OH” is practically not contained in the solution, and there is no such compound in the solid hydrate. Equations of the most important reactions:
NH 3 H 2 O (conc.) = NH 3 + H 2 O (boiling with NaOH)
NH 3 H 2 O + HCl (diluted) = NH 4 Cl + H 2 O
3(NH 3 H 2 O) (conc.) + CrCl 3 = Cr(OH) 3 ↓ + 3NH 4 Cl
8(NH 3 H 2 O) (conc.) + ZBr 2 (p) = N 2 + 6NH 4 Br + 8H 2 O (40–50 °C)
2(NH 3 H 2 O) (conc.) + 2KMnO 4 = N 2 + 2MnO 2 ↓ + 4H 2 O + 2KOH
4(NH 3 H 2 O) (conc.) + Ag2O= 2OH + 3H2O
4(NH 3 H 2 O) (conc.) + Cu(OH) 2 + (OH) 2 + 4H 2 O
6(NH 3 H 2 O) (conc.) + NiCl 2 = Cl 2 + 6H 2 O
A dilute ammonia solution (3–10%) is often called ammonia(the name was invented by alchemists), and the concentrated solution (18.5–25%) - ammonia water(produced by industry).
Related information.
Hydrogen sulfide and sulfides. Hydrogen sulfide H 2 S is a colorless gas with a pungent odor. Very poisonous, causing poisoning even at low levels in the air (about 0.01%). Hydrogen sulfide is all the more dangerous because it can accumulate in the body. It combines with the iron in hemoglobin in the blood, which can lead to fainting and death from oxygen starvation. In the presence of organic vapors, the toxicity of H 2 S increases sharply.
However, hydrogen sulfide is integral part some mineral waters (Pyatigorsk, Sernovodsk, Matsesta) used for medicinal purposes.
Hydrogen sulfide is contained in volcanic gases and is constantly formed at the bottom of the Black Sea. Hydrogen sulfide does not reach the upper layers, since at a depth of 150 m it interacts with oxygen penetrating from above and is oxidized to sulfur. Hydrogen sulfide is formed when protein rots, which is why, for example, rotten eggs smell of hydrogen sulfide.
When hydrogen sulfide is dissolved in water, weak hydrosulfide acid is formed, the salts of which are called sulfides. Sulfides of alkali and alkaline earth metals, as well as ammonium sulfide, are highly soluble in water, sulfides of other metals are insoluble and are painted in various colors, for example: ZnS - white, PbS - black, MnS - pink (Fig. 120).
Rice. 120.
Metal sulfides have different colors
Hydrogen sulfide burns. When the flame is cooled (cold objects are introduced into it), free sulfur is formed:
2H 2 S + O 2 = 2H 2 O + 2S↓.
If the flame is not cooled and excess oxygen is provided, then sulfur oxide (IV) is obtained:
2H 2 S + 3O 2 = 2H 2 O + 2SO 2.
Hydrogen sulfide is a powerful reducing agent.
Sulfur (IV) oxide, sulfurous acid and its salts. When sulfur is burned, hydrogen sulfide is completely burned and sulfides are burned, sulfur oxide (IV) SO 2 is formed, which, as noted earlier, is often also called sulfur dioxide. It is a colorless gas with a characteristic pungent odor. It exhibits typical properties of acidic oxides and is highly soluble in water, forming weak sulfurous acid. It is unstable and decomposes into its original substances:
Salts of sulfurous acid, as a dibasic acid, can be medium - sulfites, for example sodium sulfite Na 2 SO 4, and acidic - hydrosulfites, for example sodium hydrosulfite NaHSO 3. Hydrogen sulfite and sodium sulfite, like sulfur dioxide, are used to bleach wool, silk, paper and straw, and also as preservatives for preserving fresh fruits and vegetables.
Sulfuric acid and its salts. When sulfur (IV) oxide is oxidized, sulfur (VI) oxide is formed:
The reaction begins only at relatively high temperatures(420-650 °C) and occurs in the presence of a catalyst (platinum, vanadium oxides, iron, etc.).
Sulfur oxide (VI) SO 3 under normal conditions is a volatile, colorless liquid with a suffocating odor. This typical acidic oxide dissolves in water to form sulfuric acid:
H 2 O + SO 3 = H 2 SO 4.
Chemically pure sulfuric acid is a colorless, oily, heavy liquid. It has a strong hygroscopic (water-removing) property, therefore it is used for drying substances. Concentrated sulfuric acid can remove water from organic molecules, charring them. If you apply a pattern to filter paper using a solution of sulfuric acid and then heat it, the paper will turn black (Fig. 121, a) and the pattern will appear.
Rice. 121.
Carbonization of paper (a) and sugar (b) with concentrated sulfuric acid
If you place powdered sugar in a tall glass glass, moisten it with water and add concentrated sulfuric acid, stirring the contents of the glass with a glass rod, then after 1-2 minutes the contents of the glass will begin to turn black, swell and rise upward in the form of a voluminous loose mass (Fig. 121, b). The mixture in the glass becomes very hot. Reaction equation for the interaction of concentrated sulfuric acid with powdered sugar(sucrose C 12 H 22 O 11)
explains the experiment: the gases formed as a result of the reaction swell the resulting coal, pushing it out of the glass along with the stick.
Concentrated sulfuric acid dissolves sulfur oxide (VI) well; a solution of SO 3 in sulfuric acid is called oleum.
You already know the rule for diluting concentrated sulfuric acid, but let’s repeat it again: you can’t add water to the acid (why?), you should carefully pour the acid into the water in a thin stream, continuously stirring the solution.
The chemical properties of sulfuric acid largely depend on its concentration.
Dilute sulfuric acid exhibits all the characteristic properties of acids: it interacts with metals in the voltage series up to hydrogen, with the release of H2, with metal oxides (basic and amphoteric), with bases, with amphoteric hydroxides and salts.
Laboratory experiment No. 29
Properties of dilute sulfuric acid
Carry out experiments to prove that sulfuric acid exhibits the typical properties of acids.
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Since sulfuric acid is dibasic, it forms two series of salts: medium - sulfates, for example Na 2 SO 4, and acidic - hydrosulfates, for example NaHSO 4.
The reagent for sulfuric acid and its salts is barium chloride BaCl 2; sulfate ions with Ba 2+ ions form white insoluble barium sulfate, which precipitates (Fig. 122):
Rice. 122.
Qualitative reaction to sulfate ion
Concentrated sulfuric acid has very different properties from dilute acid. Thus, when H 2 SO 4 (conc) interacts with metals, hydrogen is not released. With metals to the right of hydrogen in the voltage series (copper, mercury, etc.), the reaction proceeds as follows:
The processes of oxidation and reduction that occur in this case can be written as follows:
When interacting with metals that are in the stress series before hydrogen, concentrated sulfuric acid is reduced to S, SO 2 or H 2 S, depending on the position of the metal in the stress series and the reaction conditions, for example:
Now you understand that metals in the voltage series both before and after hydrogen interact with H 2 SO 4 (conc). In this case, hydrogen is not formed, since the oxidizing agent in such a reaction is not hydrogen cations H+, as in H 2 SO 4 (dil), but sulfate ions.
Iron and aluminum are passivated with concentrated sulfuric acid, i.e. they are covered with a protective film, so concentrated acid can be transported in steel and aluminum tanks.
Being a non-volatile strong acid, concentrated sulfuric acid is capable of displacing other acids from their salts. You already know this reaction, for example the production of hydrogen chloride:
Sulfuric acid is one of the most important products used in various industries (Fig. 123). The main areas of its application: production of mineral fertilizers, metallurgy, refining of petroleum products.
Rice. 123.
Application of sulfuric acid:
1-8 - production of chemical products and goods (acids 1, explosives 2, mineral fertilizers 3, electrolytic copper 4, enamel 5, salts 6, rayon 7, medicines 8); 9 - purification of petroleum products; 10 - as an electrolyte in batteries
Sulfuric acid is also used in the production of other acids, detergents, explosives, medicines, paints, and as an electrolyte for lead batteries. Figure 124 shows the amount of sulfuric acid (in %) of total world production used in various industries.
Rice. 124.
Share of sulfuric acid consumption for various needs of industrial production
From salts of sulfuric acid highest value they already have sodium sulfate, or Glauber’s salt, Na 2 SO 4 10H 2 O, gypsum CaSO 4 2H 2 O and barium sulfate BaSO4 (where are they used?).
Copper sulfate CuSO 4 5H 2 O is used in agriculture to combat pests and plant diseases.
Sulfuric acid production. Sulfuric acid is prepared in three stages.
The chemical processes for the production of sulfuric acid can be represented as the following diagram:
1. Obtaining SO 2. Sulfur, pyrite or hydrogen sulfide are used as raw materials:
2. Obtaining SO 3. You already know this process - oxidation with oxygen is carried out using a catalyst (write down the uranium of the reaction and give its full characteristics).
3. Obtaining H 2 SO 4. But here, in contrast to the reaction you know, described by the equation:
SO 3 + H 2 O = H 2 SO 4,
The process of dissolving sulfur (VI) oxide is carried out not in water, but in concentrated sulfuric acid, which produces the familiar oleum.
The production of sulfuric acid creates a lot environmental problems. Emissions and waste from sulfuric acid plants have an extremely negative impact, causing damage to the respiratory system in humans and animals, death of vegetation and suppression of its growth, increased corrosive wear of materials, destruction of structures made of limestone and marble, acidification of soils, etc.
New words and concepts
- Hydrogen sulfide and sulfides.
- Sulfur dioxide, sulfurous acid, sulfites.
- Sulfuric acid, diluted and concentrated.
- Application of sulfuric acid.
- Sulfuric acid salts: Glauber's salt, gypsum, barium sulfate, copper sulfate.
- Production of sulfuric acid.
Tasks for independent work
- Which of the substances exhibits only reducing, only oxidizing, or both oxidizing and reducing properties: sulfur, hydrogen sulfide, sulfur oxide (IV), sulfuric acid? Why? Support your answer with equations for the corresponding reactions.
- Describe: a) sulfur dioxide; b) sulfur oxide (VI) according to plan: preparation, properties, application. Write the equations for the corresponding reactions.
- Write reaction equations characterizing the properties of dilute sulfuric acid as an electrolyte. Which property is a redox process? What reactions can be classified as ion exchange reactions? Consider them from the point of view of the theory of electrolytic dissociation.
- Write the equations for the reactions underlying the production of sulfuric acid according to the diagram given in the paragraph.
- 40 g of sulfur oxide (VI) (no.) was dissolved in 400 ml of water. Calculate the mass fraction of sulfuric acid in the resulting solution.
- Characterize the reaction for the synthesis of sulfur (VI) oxide using all the reaction classifications you have studied.
- 500 g of copper sulfate was dissolved in 5 liters of water. Calculate the mass fraction of copper (II) sulfate in the resulting solution.
- Why is sulfuric acid called the “bread of the chemical industry”?
- (hydrogen sulfide) H2S, a colorless gas with the smell of rotten eggs; melting point?85.54.C, boiling point?60.35.C; at 0.C it liquefies under a pressure of 1 MPa. Reducing agent. A by-product during the refining of petroleum products, coking of coal, etc.; formed during decomposition... ... Big Encyclopedic Dictionary
HYDROGEN Sulfide- (H2S), a colorless, poisonous gas with the smell of rotten eggs. Formed during decay processes, found in crude oil. Obtained by the action of sulfuric acid on metal sulfides. Used in traditional QUALITATIVE ANALYSIS. Properties: temperature... ... Scientific and technical encyclopedic dictionary
HYDROGEN Sulfide- HYDROGEN Sulfide, hydrogen sulfide, many others. no, husband (chem.). A gas produced by the rotting of protein substances, giving off the smell of rotten eggs. Dictionary Ushakova. D.N. Ushakov. 1935 1940 ... Ushakov's Explanatory Dictionary
HYDROGEN Sulfide- HYDROGEN SULFIDE, huh, husband. A colorless gas with a sharp, unpleasant odor, formed during the decomposition of protein substances. | adj. hydrogen sulfide, oh, oh. Ozhegov's explanatory dictionary. S.I. Ozhegov, N.Yu. Shvedova. 1949 1992 … Ozhegov's Explanatory Dictionary
hydrogen sulfide- noun, number of synonyms: 1 gas (55) ASIS Dictionary of Synonyms. V.N. Trishin. 2013… Synonym dictionary
HYDROGEN Sulfide- colorless poisonous gas H2S with an unpleasant specific odor. It has slightly acidic properties. 1 liter of C. at t 0 °C and a pressure of 760 mm is 1.539 g. It is found in oils, natural waters, and gases of biochemical origin, such as... ... Geological encyclopedia
HYDROGEN Sulfide- HYDROGEN Sulfide, H2S (molecular weight 34.07), a colorless gas with a characteristic odor of rotten eggs. A liter of gas under normal conditions (0°, 760 mm) weighs 1.5392 g. Boiling temperature 62°, melting 83°; S. is part of gaseous emissions... ... Great Medical Encyclopedia
hydrogen sulfide- - Topics of biotechnology EN hydrogen sulfide ... Technical Translator's Guide
hydrogen sulfide- HYDROGEN SULFIDE, a, m A colorless gas with a sharp, unpleasant odor, formed during the decomposition of protein substances and representing a compound of sulfur with hydrogen. Hydrogen sulfide is found in some mineral waters and medicinal mud and is used... ... Explanatory dictionary of Russian nouns
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