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Sulfur is a fairly common chemical element in nature (the sixteenth most abundant in earth's crust and the sixth - in natural waters). Both native sulfur (the free state of the element) and its compounds are found.

Sulfur in nature

Among the most important natural ones are iron pyrite, sphalerite, galena, cinnabar, and stibnite. The World Ocean contains mainly magnesium and sodium, which cause the hardness of natural waters.

How is sulfur obtained?

Sulfur ores are mined using different methods. The main way to obtain sulfur is to smelt it directly at its deposits.

Open-pit mining involves the use of excavators to remove rock layers that cover the sulfur ore. After the ore layers are crushed by explosions, they are sent to a sulfur smelter.

In industry, sulfur is obtained as a by-product of processes in smelting furnaces and oil refining. It is present in large quantities in natural gas(in the form of sulfur dioxide or hydrogen sulfide), during the extraction of which it is deposited on the walls of the equipment used. Fine sulfur captured from gas is used in chemical industry as raw materials for the production of various products.

This substance can also be obtained from natural sulfur dioxide. For this, the Klaus method is used. It involves the use of “sulfur pits” in which sulfur is degassed. The result is modified sulfur, which is widely used in asphalt production.

Main allotropic modifications of sulfur

Sulfur is characterized by allotropy. Known a large number of allotropic modifications. The most famous are orthorhombic (crystalline), monoclinic (needle-shaped) and plastic sulfur. The first two modifications are stable, the third, upon hardening, turns into a rhombic one.

Physical properties characterizing sulfur

Molecules of the orthorhombic (α-S) and monoclinic (β-S) modifications each contain 8 sulfur atoms, which are connected in a closed cycle by single covalent bonds.

Under normal conditions, sulfur has an orthorhombic modification. It is a yellow crystalline solid with a density of 2.07 g/cm 3 . Melts at 113 °C. The density of monoclinic sulfur is 1.96 g/cm 3, its melting point is 119.3 °C.

When melted, sulfur expands in volume and becomes a yellow liquid, which turns brown at a temperature of 160 °C and turns into a viscous dark brown mass when it reaches about 190 °C. At temperatures above this value, the viscosity of sulfur decreases. At about 300 °C it again turns into a liquid fluid state. This is explained by the fact that during the heating process, sulfur polymerizes, increasing the length of the chain with increasing temperature. And when the temperature reaches above 190 °C, destruction of the polymer links is observed.

When molten sulfur is cooled naturally in cylindrical crucibles, so-called lump sulfur is formed - orthorhombic crystals large sizes, having a distorted shape in the form of octahedra with partially “cut off” faces or corners.

If the molten substance is subjected to rapid cooling (for example, using cold water), then you can get plastic sulfur, which is an elastic rubber-like mass of brownish or dark red color with a density of 2.046 g/cm 3 . This modification, unlike the orthorhombic and monoclinic ones, is unstable. Gradually (over several hours) it changes color to yellow, becomes fragile and turns into a rhombic one.

When sulfur vapor (highly heated) is frozen with liquid nitrogen, its purple modification is formed, which is stable at temperatures below minus 80 °C.

Sulfur is practically insoluble in an aqueous environment. However, it is characterized by good solubility in organic solvents. Conducts electricity and heat poorly.

The boiling point of sulfur is 444.6 °C. The boiling process is accompanied by the release of orange-yellow vapors, consisting mainly of S 8 molecules, which dissociate upon subsequent heating, resulting in the formation of equilibrium forms S 6, S 4 and S 2. Further, when heated, large molecules decompose, and at temperatures above 900 degrees, the vapors consist almost exclusively of S 2 molecules, which dissociate into atoms at 1500 ° C.

What chemical properties does sulfur have?

Sulfur is a typical non-metal. Chemically active. Oxidatively - The reducing properties of sulfur are manifested in relation to many elements. When heated, it easily combines with almost all elements, which explains its essential presence in metal ores. The exceptions are Pt, Au, I 2, N 2 and inert gases. The oxidation states that sulfur exhibits in compounds are -2, +4, +6.

The properties of sulfur and oxygen cause it to burn in air. The result of this interaction is the formation of sulfurous (SO 2) and sulfuric (SO 3) anhydrides, which are used to produce sulfurous and sulfuric acids.

At room temperature The reducing properties of sulfur appear only in relation to fluorine, in reaction with which it is formed:

• S + 3F 2 = SF 6.

When heated (in the form of a melt), it interacts with chlorine, phosphorus, silicon, and carbon. As a result of reactions with hydrogen, in addition to hydrogen sulfide, it forms sulfanes, combined general formula H 2 S X.

The oxidizing properties of sulfur are observed when interacting with metals. In some cases, quite violent reactions can be observed. As a result of interaction with metals, compounds) and polysulfides (polysulfur metals) are formed.

When heated for a long time, it reacts with concentrated oxidizing acids, thereby oxidizing.

Sulfur dioxide

Sulfur (IV) oxide, also called sulfur dioxide and sulfur dioxide, is a gas (colorless) with a pungent, asphyxiating odor. It has the property of liquefying under pressure at room temperature. SO 2 is an acidic oxide. Characterized by good solubility in water. This produces a weak, unstable sulfurous acid, which exists only in an aqueous solution. As a result of the interaction of sulfur dioxide with alkalis, sulfites are formed.

It has quite high chemical activity. The most pronounced are restorative Chemical properties sulfur(IV) oxide. Such reactions are accompanied by an increase in the degree of sulfur oxidation.

The oxidative chemical properties of sulfur oxide occur in the presence of strong reducing agents (for example, carbon monoxide).

Sulfur trioxide

Sulfur trioxide (sulfuric anhydride) - sulfur (VI). Under normal conditions, it is a colorless, highly volatile liquid characterized by a suffocating odor. It tends to harden at temperatures below 16.9 degrees. In this case, a mixture of different crystalline modifications of solid sulfur trioxide is formed. The high hygroscopic properties of sulfur oxide cause it to “smoke” in humid air conditions. As a result, droplets of sulfuric acid are formed.

Hydrogen sulfide

Hydrogen sulfide is a binary chemical compound of hydrogen and sulfur. H 2 S is a poisonous colorless gas, characteristic features which has a sweetish taste and smell of rotten eggs. Melts at minus 86 °C, boils at minus 60 °C. Thermally unstable. At temperatures above 400 °C, hydrogen sulfide decomposes into S and H2. Characterized by good solubility in ethanol. It dissolves poorly in water. As a result of dissolution in water, weak hydrosulfide acid is formed. Hydrogen sulfide is a strong reducing agent.

Flammable. When it burns in the air, you can see a blue flame. In high concentrations it can react with many metals.

Sulfuric acid

Sulfuric acid (H 2 SO 4) can be of different concentrations and purities. In an anhydrous state, it is a colorless, odorless, oily liquid.

The temperature at which the substance melts is 10 °C. The boiling point is 296 °C. It dissolves well in water. When sulfuric acid dissolves, hydrates are formed and a large amount of heat is released. The boiling point of all aqueous solutions at a pressure of 760 mm Hg. Art. exceeds 100 °C. An increase in boiling point occurs with increasing acid concentration.

The acidic properties of a substance appear when it interacts with bases. H 2 SO 4 is a dibasic acid, due to which it can form both sulfates (medium salts) and hydrosulfates ( acid salts), most of which are soluble in water.

The properties of sulfuric acid are most clearly manifested in redox reactions. This is explained by the fact that in the composition of H 2 SO 4 sulfur has the highest oxidation state (+6). An example of the manifestation of the oxidizing properties of sulfuric acid is the reaction with copper:

• Cu + 2H 2 SO 4 = CuSO 4 + 2H 2 O + SO 2.

Sulfur: beneficial properties

Sulfur is a trace element essential for living organisms. Is integral part amino acids (methionine and cysteine), enzymes and vitamins. This element takes part in the formation of the tertiary structure of the protein. The amount of chemically bound sulfur contained in proteins ranges from 0.8 to 2.4% by weight. The content of the element in the human body is about 2 grams per 1 kg of weight (that is, approximately 0.2% is sulfur).

The beneficial properties of the microelement are difficult to overestimate. By protecting the protoplasm of the blood, sulfur is an active assistant to the body in the fight against harmful bacteria. Blood clotting depends on its quantity, that is, the element helps maintain its sufficient level. Sulfur also plays an important role in maintaining normal values the concentration of bile produced by the body.

It is often called the “beauty mineral” because it is essential for maintaining healthy skin, nails and hair. Sulfur has the ability to protect the body from various types negative impact environment. This helps slow down the aging process. Sulfur cleanses the body of toxins and protects against radiation, which is especially important nowadays, given the current environmental situation.

An insufficient amount of microelement in the body can lead to poor elimination of toxins, decreased immunity and vitality.

Sulfur is a participant in bacterial photosynthesis. It is a component of bacteriochlorophyll, and hydrogen sulfide is a source of hydrogen.

Sulfur: properties and industrial applications

Sulfur is most widely used for. Also, the properties of this substance allow it to be used for the vulcanization of rubber, as a fungicide in agriculture and even medicinal product(colloidal sulfur). In addition, sulfur is used for the production of matches and is included in sulfur-bitumen compositions for the production of sulfur asphalt.

1.1. Historical reference

Sulfur is one of the few substances that has been known since ancient times; it was used by the first chemists. One of the reasons for the popularity of sulfur is the prevalence of native sulfur in the countries of ancient civilizations. It was developed by the Greeks and Romans, and sulfur production increased significantly after the invention of gunpowder.

1.2. Place of sulfur in the Periodic Table chemical elements Mendeleev

Sulfur is located in group 16 of Mendeleev's Periodic Table of Chemical Elements.

The outer energy level of the sulfur atom contains 6 electrons, which have the electronic configuration 3s 2 3p 4. In compounds with metals, sulfur exhibits a negative oxidation state of elements -2, in compounds with oxygen and other active non-metals - positive +2, +4, +6. Sulfur is a typical non-metal; depending on the type of transformation, it can be an oxidizing agent and a reducing agent.

1.3. Prevalence in nature

Sulfur is quite widespread in nature. Its content in the earth's crust is 0.0048%. A significant part of sulfur is found in the native state.

Sulfur is also found in the form of sulfides: pyrite, chalcopyrite and sulfates: gypsum, celestine and barite.

Many sulfur compounds are found in oil (thiophene C 4 H 4 S, organic sulfides) and petroleum gases (hydrogen sulfide).

1.4. Allotropic modifications of sulfur

The existence of allotropic modifications of sulfur is associated with its ability to form stable homochains – S – S –. The stability of the chains is explained by the fact that the bonds – S – S – are stronger than the bond in the S 2 molecule. Sulfur homochains have a zigzag shape, since electrons from mutually perpendicular p-orbitals take part in their formation.

There are three allotropic modifications of sulfur: orthorhombic, monoclinic and plastic. The rhombic and monoclinic modifications are constructed from cyclic S8 molecules located at the sites of the rhombic and monoclinic lattices.

The S8 molecule has the shape of a crown, the lengths of all bonds – S – S – are equal to 0.206 nm and the angles are close to tetrahedral 108°.

In rhombic sulfur, the smallest elementary volume has the shape of a rectangular parallelepiped, and in the case of monoclinic sulfur, the elementary volume is allocated in the form of a beveled parallelepiped.

Orthorhombic sulfur crystal Monoclinic sulfur crystal

The plastic modification of sulfur is formed by helical chains of sulfur atoms with left and right axes of rotation. These chains are twisted and pulled in one direction.

Orthorhombic sulfur is stable at room temperature. When heated, it melts, turning into a yellow, easily mobile liquid; with further heating, the liquid thickens, as long polymer chains are formed in it. When the melt is slowly cooled, dark yellow needle-shaped crystals of monoclinic sulfur are formed, and if you pour molten sulfur into cold water, you get plastic sulfur - a rubber-like structure consisting of polymer chains. Plastic and monoclinic sulfur are unstable and spontaneously transform into orthorhombic sulfur.

1.5. Physical properties of sulfur

Sulfur is a solid, brittle, yellow substance, practically insoluble in water, not wetted by water and floats on its surface. It dissolves well in carbon disulfide and other organic solvents, and is a poor conductor of heat and electric current. When melted, sulfur forms a readily mobile yellow liquid, which darkens at 160°C, its viscosity increases, and at 200°C sulfur becomes dark brown and viscous, like resin. This is explained by the destruction of ring molecules and the formation of polymer chains. Further heating causes the chains to break, and the liquid sulfur again becomes more mobile. Sulfur vapors range in color from orange-yellow to straw-yellow. Steam consists of molecules of the composition S 8, S 6, S 4, S 2. At temperatures above 150 °C, the S2 molecule dissociates into atoms.

The physical properties of allotropic modifications of sulfur are given in the table:

Sulfur is one of the substances known to mankind from time immemorial. Even the ancient Greeks and Romans found various practical applications for it. Pieces of native sulfur were used to perform the ritual of expelling evil spirits. So, according to legend, Odysseus, returning to his home after long wanderings, first ordered it to be fumigated with sulfur. There are many references to this substance in the Bible.

In the Middle Ages, sulfur occupied an important place in the arsenal of alchemists. They believed that all metals consist of mercury and sulfur: the less sulfur, the more noble the metal. Practicalinterest in this substance in Europe has increased in XIII - XIV centuries, after the advent of gunpowder and firearms.

Sulfur mining. Engraving from the book “On Mining and Metallurgy” by G. Agricola. Edition 1557. Ore containing sulfur is heated in wide clay pots A with long beak-shaped spouts lowered into special holes in the receiver B, closed with a lid. C. Molten sulfur is scooped out of the receiver using ladles and poured into forms.

Rhombic sulfur.

A monoclinic modification of sulfur crystallizes from the melt.

Plastic sulfur is elastic, like rubber.

The main supplier of sulfur was Italy.

Today, sulfur is used as a raw material for the production of sulfuric acid, with vulcanization rubber, in organic synthesis. Sulfur powder is used in medicine as an external disinfectant.

Sulfur forms several allotropic modifications. Stable at room temperature rhombic series It is a yellow powder, insoluble in water. Upon crystallization from chloroform CHC l 3 or from carbon disulfide CS 2 it stands out in the form of transparent crystals of octahedral shape. Orthorhombic sulfur consists of cyclic molecules S 8, shaped like a crown. At 113 °C it melts, turning into a yellow, easily mobile liquid. With further heating, the melt thickens, as long polymer chains are formed in it. And if you heat sulfur to 445 °C, it boils. By pouring boiling sulfur in a thin stream into cold water, you can get plastic sulfur- rubber-like modification consisting of poly-

measuring chains. As the melt cools slowly, dark yellow needle-shaped crystals form monoclinic sulfur(t pl =119°С). Like rhombic sulfur, this modification consists of molecules S 8 . At room temperature, plastic and monoclinic sulfur are unstable and spontaneously transform into orthorhombic sulfur powder.

When heated, sulfur reacts with many metals (iron, aluminum, mercury) and non-metals (oxygen, halogens, hydrogen). “The nature of sulfur is fiery, combustible... [Sulfur] burns completely, evaporating into smoke,” is written in one alchemical treatise. Indeed, when sulfur burns in air or oxygen, sulfur oxide is formed ( IV ), or sulfur gas, SO2, containing an admixture (about 3% by volume) of higher sulfur oxide, or sulfuric anhydride, SO 3,. Sulfur SO2 gas - a colorless gas with a suffocating, pungent odor. When dissolved in water (at 0 °C, 1 volume of water dissolves more than 70 volumes SO 2) sulfurous acid is formed H2SO3, which is known only in solutions. However, its salts are sulfites (for example, Na 2 SO 3) and hydrosulfites(NaHSO 3) - can easily be isolated in solid form.

In laboratory conditions to obtain SO, act on solid sodium sulfite with concentrated sulfuric acid:Na 2 SO 3 +2H 2 SO 4 =2NaHSO 4+ SO 2 - + H2O.

Combustion of sulfur in oxygen.

Molecule structure SO 2

Structure of sulfur:

1—orthorhombic and monoclinic, S 8;

2— plastic, S n

*Vulcanization is the process of converting rubber into rubber through the formation of sulfide “bridges” between individual polymer molecules. The resulting cross-linked polymer has a spatial structure and is characterized by increased mechanical strength.

In industry SO 2 obtained by roasting sulfide ores, for example pyrite: 4FeS 2 +11 O 2 = 2Fe 2 O 3 + 8SO 2, or when burning sulfur. Sulfur dioxide is an intermediate product in the production of sulfuric acid. It is also used (together with sodium hydrosulfites NaHSO 3 and calcium Ca(HSO 3) 2) for extracting cellulose from wood. Trees and shrubs are fumigated with this gas to kill agricultural pests.

Sulfuric anhydride SO 3 at room temperature it is a colorless, easily volatile liquid (t bale =45 °C), which over time turns into an asbestos-like modification consisting of shiny silky crystals. Sulfuric anhydride fibers are stable only in a sealed container. Absorbing air moisture, they turn into a thick, colorless liquid - oleum (from lat. oleum - "oil"). Although formally oleum can be considered a solution SO 3 in H 2 SO 4, it is actually a mixture of various pyrosulfuric acids: H 2 S 2 O 7, H 2 S 3 O 10, etc. With water SO 3 interacts very energetically: in this case, so much heat is released that the tiny droplets of sulfuric acid formed create fog. You need to work with this substance with extreme caution.

The structure of the asbestos-like modification of sulfuric anhydride.

Asbestos modification SO3.

Sulfur dioxide exhibits a strong bleaching effect: a red rose fluffed into a flask with SO2, loses its color.

Sulfuric acid H2SO4 - a heavy oily, colorless liquid, miscible with water in any proportions. At 10 °C it hardens, forming a transparent glassy mass. When heated, 100 percent sulfuric acid easily loses sulfuric anhydride until its concentration is 98 %. It is this acid that is usually used in laboratories (concentrated sulfuric acid,t kip =338 °C).

Remember that you need to pour the acid into the water in a thin stream with constant stirring. Under no circumstances should you pour water into sour water! Due to strong heating, the water will boil, and hot splashes of sulfuric acid solution can get into your face.

Dilute sulfuric acid exhibits all the properties inorganic acids: reacts with basic oxides, bases and active metals to release hydrogen. H2SO4 refers to strong acids, in an aqueous solution of the acid its molecules

Structure (SO 3) 3 .

*Asbestos-like SO 3 (t pl = 32 °C) is a crystalline polysulfuric acid consisting of long chains HO - (S(O) 2 - O - ) n - OH. However, in fact this is pure sulfuric anhydride, since the length of such a chain is ( n) is 10 5 i.e., there are two hydrogen atoms for every 10 5 sulfur atoms. Liquid modification of sulfuric anhydride at room temperature(t pl = 17 °C), consists of cyclic molecules(SO 3) 3 .

does not exist: they decay into hydrogen ions and hydrogen sulfate ions(HSO - 4), which dissociate only when highly diluted.

Concentrated sulfuric acid is a strong oxidizing agent. It reacts both with active metals and with those standing in the voltage series to the right of hydrogen - copper, silver, mercury. Metal oxidizes and sulfuric acid is reduced before sulfur, hydrogen sulfide (in reaction with zinc, magnesium) or to sulfur dioxide, as happens when interacting with an inactive metal - copper:

Cu+ 2H 2 SO 4 = CuSO 4 + SO 2 - +2H 2 O.

Strong (50-70 percent) sulfuric acid easily oxidizes iron:

2Fe+6H 2 SO 4 =Fe 2 (SO 4) 3 + 3SO 2 - +6H 2 O.

At the same time, in the cold, oleum does not react with iron and aluminum.

Concentrated sulfuric acid is capable of charring many organic substances (sugar, paper, cotton wool). In case of accidental hit H2SO4 If applied to the skin, you must immediately rinse it off with a stream of water, and then treat the burn site with a weak soda solution.

Mention of sulfuric acid was first found among Arab and European alchemists. It was obtained by calcination in air inkstone(vitriol, or hydrated ferrous sulfate ( II), FeSO 4 .7H 2 O):

2FeSO 4 = Fe 2 O 3 + SO 3 - + SO 2 - or a mixture of sulfur and saltpeter: 6KNO 3 + 5S = 3K 2 SO 4 + 2SO 3 - + 3N 2 - , and the released vapors of sulfuric anhydride condensed. Absorbing moisture, they turned into oleum. Depending on the cooking method H2SO4 called oil of vitriol(oleum vitrioli) or sulfur oil(oleum sulfuris). In 1595, the alchemist Andreas Libavius ​​(1550-1616) established the identity of both substances.

For a long time, oil of vitriol was not found wide application. Interest in it increased greatly after XVIII V. The process of obtaining indigo carmine, a stable blue dye, from indigo was discovered. The first factory for the production of sulfuric acid was founded near London in 1736. The process was carried out in lead chambers, at the bottom of which water was poured. In the upper part of the chamber, a molten mixture of saltpeter and sulfur was burned, then air was introduced into it. The procedure was repeated until an acid of the required concentration was formed at the bottom of the container. In this case, the following chemical transformations occurred:

S +O 2 = SO 2 2KNO 3 +S = K 2 SO 4 +2NO

2NO + O 2 = 2NO 2 NO 2 + SO 2 + H 2 O = H 2 SO 4 + NO.

In the XIX V. Has the method been improved: instead of saltpeter, they began to use nitric acid(when decomposed in the chamber it gives NO 2). To return nitrous gases to the system, special towers were constructed, which gave the name to the whole process - tower process. Factories operating using the tower method still exist today.

However, now for the production of sulfuric acid it is mainly used contact method developed in 1831. Using this method, oxidation SO 2 to SO 3

(2SO 2 +O 2 « 2SO 3) carried out on a catalyst - vanadium oxide ( V ) V 2 O 5 .

Ser In special installations, the anhydride is absorbed by concentrated sulfuric acid. This produces oleum. It is stored in iron tanks and, as necessary, converted to sulfuric acid.



Installation for producing sulfuric acid by burning sulfur in the presence of nitrate.

Middle XVIII V. A mixture of sulfur and saltpeter prepared in advance is loaded into a furnace (1), heated by coals.

The resulting gases reach glass vessel(2), where they interact with water vapor. The resulting oleum is collected in flasks (3).

Mineral pyrite FeS 2 is an iron disulfide, a salt of a weak acid H2S2, constructed similarly to hydrogen peroxide H 2 O 2 .

a) Pour up to half of its volume pieces of cutting sulfur into a test tube fixed in a holder and heat it very carefully, shaking all the time. The sulfur will begin to melt, forming a yellow, mobile liquid. Above 160° the liquid darkens, and at 200° it becomes dark brown and so viscous that it does not pour out of the test tube. Above 250°, the viscosity decreases again and at 400°, sulfur turns into a highly mobile dark brown liquid, which boils at 444.5°, forming orange-yellow vapors. Explain the changes that occur when molten sulfur is heated.

b) Pour boiling sulfur in a thin stream into a glass with cold water. If the sulfur flares up, after pouring it out, cover the opening of the test tube with a crucible lid or a piece of asbestos. Remove the resulting mass from the water and make sure it is plastic. Save the resulting plastic sulfur in order to trace the transition from amorphous to crystalline form.

Preparation of rhombic sulfur

Place 2-3 pea-sized pieces of cutting sulfur in a test tube, add 2 ml of carbon disulfide and, shaking, dissolve the sulfur. Carbon disulfide is a flammable liquid, and all work with it must be carried out away from fire. Pour a few drops of the resulting solution onto a watch glass. Allow the carbon disulfide to evaporate and observe the release of orthorhombic sulfur crystals.

3) Sublimation of red phosphorus

Put a little red phosphorus in a test tube, cover it with cotton wool and, fixing it horizontally in a stand, lightly heat it with a burner flame. The red phosphorus evaporates and a white coating deposits on the cold parts of the test tube. Perform the experiment carefully, always making sure that the phosphorus vapor does not ignite when leaving the test tube.

4) Phosphorus combustion under water

a) Place a piece of white phosphorus in a glass of water and heat the water to 60-70°. Then pass a weak current of oxygen through the gasometer, holding the outlet tube so that it touches the phosphorus. The last one lights up. Write the reaction equation.

b) Do the same, replacing white phosphorus with red. Red phosphorus does not burn under water.

1. Phosphoric anhydride and its properties

(Work is done under traction)

Place 0.4-0.5 g of red phosphorus in a porcelain cup (or on the crucible lid), placed on an asbestos mesh. Place a dry funnel above the cup at a short distance (about 0.5 cm) from the mesh. Light the phosphorus with a hot glass rod. Phosphoric anhydride, formed during the combustion of phosphorus, is deposited on the walls of the funnel in the form of a white crystalline mass similar to cigar.

When all the phosphorus has burned, place the funnel in the tripod ring and leave for a while. Phosphoric anhydride spreads very quickly. What property of P 2 O 5 does this phenomenon indicate?

2. Preparation of phosphoric acids

a) Wash off the phosphoric anhydride obtained in the previous experiment with distilled water from the walls of the funnel into a test tube. When the solution becomes transparent, pour it a little into another test tube and add an excess of AgN0 3 solution to the last one. A white precipitate of AgP0 3 is formed. Write the reaction equations.

b) Pour the rest of the H 3 PO 4 solution into a glass, add 10-15 ml of water and 1-2 ml of concentrated HNO 3

3. Action alkali metals to the water

(The work is done behind glass fume hood!)

Take three porcelain cups with water. Cut off a small piece of lithium, sodium and potassium with a knife, dry them with filter paper, throw them into water: lithium in one cup, sodium in another, potassium in a third. Observe the reaction progress through the glass of a fume hood. Glass protection is necessary due to splashing that occurs at the end of reactions. Note that potassium reacts most energetically with water, and lithium reacts least energetically. Test the resulting solutions with litmus or phenolphthalein. Write reaction equations.

Obtaining potassium nitrate

Add 7.5 g of KS1 to a glass containing 20 ml of water and dissolve with heating; then add 8.5 g of crushed NaN03. Boil the contents of the glass for several minutes, then quickly filter the liquid from the formed NaCl precipitate using a shortened glass funnel (with the tube cut off). Allow the solution to cool and observe the precipitation of KN0 3 crystals. Separate the crystals by decanting the mother liquor and dry them between sheets of filter paper. Explain the phenomena observed during the experiment based on the solubility of salts that can form in the solution

5. Reaction of opening of Na - and K ions

a) Pour a neutral solution of some sodium salt into a test tube and add cold a strong (preferably freshly prepared) solution of acidic potassium pyroantimony K 2 H 2 Sb 2 0 7 . Observe the precipitation of a white crystalline precipitate of Na 2 H 2 Sb 2 0 7. If necessary, precipitation can be accelerated by rubbing a glass rod against the walls of the test tube. Write the reaction equation in molecular and ionic forms.

b) Add a solution of acidic sodium tartrate NaHC4H 4 0 6 to a neutral solution of some potassium salt and shake. Observe the precipitation of a white crystalline precipitate of KHC4H4O6. Write the reaction equation in molecular and ionic forms.

Questions

1. What is common in the processes of sodium combustion in chlorine, the interaction of sodium with water and the interaction of sodium with sulfuric acid?

2. Which of the following potassium salts will undergo noticeable hydrolysis: KC1, KNO3, K2S, K2CO3, CH3COOC?

3. Why is plastic sulfur not completely dissolved in carbon disulfide?

4. How can you remove hydrogen sulfide from a mixture of gases?

5. Is it possible to use nitric acid to obtain hydrogen sulfide from its salts?

6. Why is hydrogen sulfide a reducing agent and does not exhibit oxidizing properties?