Introduction

Phosphorus (lat. Phosphorus) P - a chemical element of group V periodic system Mendeleev atomic number 15, atomic mass 30.973762(4). Consider the structure of the phosphorus atom. There are five electrons in the outer energy level of the phosphorus atom. Graphically it looks like this:

1s 2 2s 2 2p 6 3s 2 3p 3 3d 0

In 1699 the Hamburg alchemist X. Brand, in search of a "philosopher's stone", supposedly capable of turning base metals into gold, when evaporating urine with coal and sand, isolated a white, waxy substance that could glow.

The name "phosphorus" comes from the Greek. "phos" - light and "phoros" - carrier. In Russia, the term "phosphorus" was introduced in 1746 by M.V. Lomonosov.

The main compounds of phosphorus include oxides, acids and their salts (phosphates, dihydrophosphates, hydrophosphates, phosphides, phosphites).

A lot of substances containing phosphorus are found in fertilizers. Such fertilizers are called phosphate fertilizers.

Phosphorus as an element and as a simple substance

Phosphorus in nature

Phosphorus is one of the common elements. The total content in the earth's crust is about 0.08%. Due to its easy oxidizability, phosphorus occurs in nature only in the form of compounds. The main minerals of phosphorus are phosphorites and apatites, of the latter, fluorapatite 3Ca 3 (PO 4) 2 * CaF 2 is the most common. Phosphorites are widely distributed in the Urals, the Volga region, Siberia, Kazakhstan, Estonia, Belarus. The largest deposits of apatite are located on the Kola Peninsula.

Phosphorus is an essential element of living organisms. It is present in bones, muscles, brain tissue and nerves. Made from phosphorus ATP molecules- adenosine triphosphoric acid (ATP - collector and carrier of energy). The body of an adult contains on average about 4.5 kg of phosphorus, mainly in combination with calcium.

Phosphorus is also found in plants.

Natural phosphorus consists of only one stable isotope, 31 P. Today, six radioactive isotopes of phosphorus are known.

Physical properties

Phosphorus has several allotropic modifications - white, red, black, brown, violet phosphorus, etc. The first three of these are the most studied.

White phosphorus- a colorless, yellowish crystalline substance that glows in the dark. Its density is 1.83 g/cm3. Insoluble in water, soluble in carbon disulfide. It has a characteristic garlic odor. Melting point 44°C, self-ignition temperature 40°C. To protect white phosphorus from oxidation, it is stored under water in the dark (there is a transformation into red phosphorus in the light). In the cold, white phosphorus is brittle, at temperatures above 15°C it becomes soft and can be cut with a knife.

Molecules of white phosphorus have a crystal lattice, in the nodes of which there are P 4 molecules, which have the shape of a tetrahedron.

Each phosphorus atom is connected by three?-bonds with the other three atoms.

White phosphorus is poisonous and gives difficult-to-heal burns.

red phosphorus- powdery substance of dark red color, odorless, does not dissolve in water and carbon disulfide, does not glow. Ignition temperature 260°C, density 2.3 g/cm 3 . Red phosphorus is a mixture of several allotropic modifications that differ in color (from scarlet to purple). The properties of red phosphorus depend on the conditions for its preparation. Not poisonous.

black phosphorus similar in appearance to graphite, greasy to the touch, has semiconductor properties. Density 2.7 g/cm 3 .

Red and black phosphorus have an atomic crystal lattice.

Chemical properties

Phosphorus is a non-metal. In compounds, it usually exhibits an oxidation state of +5, less often - +3 and -3 (only in phosphides).

Reactions with white phosphorus are easier than with red.

I. Interaction with simple substances.

1. Interaction with halogens:

2P + 3Cl 2 = 2PCl 3 (phosphorus (III) chloride),

PCl 3 + Cl 2 = PCl 5 (phosphorus (V) chloride).

2. Interaction with non-metals:

2P + 3S = P 2 S 3 (phosphorus (III) sulfide.

3. Interaction with metals:

2P + 3Ca = Ca 3 P 2 (calcium phosphide).

4. Interaction with oxygen:

4P + 5O 2 = 2P 2 O 5 (phosphorus (V) oxide, phosphoric anhydride).

II. Interaction with complex substances.

3P + 5HNO 3 + 2H 2 O \u003d 3H 3 PO 4 + 5NO ^.

Receipt

Phosphorus is obtained from crushed phosphorites and apatites, the latter are mixed with coal and sand and calcined in furnaces at 1500 ° C:

2Ca 3 (PO 4) 2 + 10C + 6SiO 2 6CaSiO 3 + P 4 ^ + 10CO ^.

Phosphorus is released in the form of vapors, which condense in the receiver under water, forming white phosphorus.

When heated to 250-300°C in the absence of air, white phosphorus turns red.

Black phosphorus is obtained by prolonged heating of white phosphorus at very high pressure (200°C and 1200 MPa).

Application

Red phosphorus is used in the manufacture of matches (see figure). It is part of the mixture applied to side surface matchbox. The main component of the composition of the match head is Bertolet's salt KClO 3 . From the friction of the match head on the spread, the phosphorus particles ignite in air. As a result of the oxidation reaction of phosphorus, heat is released, leading to the decomposition of Berthollet salt.

The resulting oxygen contributes to the ignition of the match head.

Phosphorus is used in metallurgy. It is used to obtain conductors and is part of some metallic materials, such as tin bronzes.

Phosphorus is also used in the production of phosphoric acid and pesticides (dichlorvos, chlorophos, etc.).

White phosphorus is used to create smoke screens, since it produces white smoke when it burns.

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OBTAINING WHITE PHOSPHORUS

When conducting experiments, it must be taken into account that white phosphorus and its vapors are poisonous; in contact with the skin, it leaves painful and long-term healing wounds ( see white phosphorus regulations).

Experience. Obtaining phosphorus as a result of the interaction of calcium orthophosphate, coal and silicon dioxide.

The reaction proceeds according to the equation:

Ca 3 (PO 4) 2 + 5C + 3SiO 2 \u003d 2P + 3CaSiO 3 + 5CO -282 kcal.

This experience makes it possible to obtain white and red phosphorus and observe its cold flame.

The reaction chamber is a refractory glass flask with a capacity of 2 l with two tubes. Flask diameter 150 mm, the length of the tubes is about 50 mm, inner diameter 40 mm.

When assembling the apparatus, the flask is mounted, as shown in the figure, on a tripod ring wrapped in asbestos and fixed at the top in the tripod clamp. Both tubes are closed with rubber stoppers, in the middle of which there is one hole for carbon electrodes and one hole for gas inlet and outlet on the side. Bottom electrode with a diameter of about 12 mm insert so that its end does not reach the middle of the flask. At the end of the electrode inserted into the flask, a small iron sleeve is fixed, which should be a support for a ceramic crucible with a hole at the bottom. The coupling used must have a screw thread and a brass screw; coupling diameter approx. 9 mm. The sleeve is screwed on so that one side of the sleeve is above the end of the electrode. A ceramic crucible (with an upper diameter of less than 40 mm), into the hole in the bottom of which the tip of the electrode is inserted. A copper sleeve is attached to the lower end of the electrode, which serves to connect the electrode to the electrical wire.

A thick-walled glass refractory tube about 100 cm long is inserted into the cork of the upper tube. ml in such a way that it is approximately 10 mm entered the flask. The upper carbon electrode, which may be thinner than the lower one, should easily pass through this tube. On the upper end of the glass tube (with melted edges) and the electrode passing through it, put on a piece of rubber tube 50 mm. The upper electrode is strengthened in such a way that its pointed end is at a distance of 8-10 mm from the top end of the bottom electrode. At the upper end of the upper electrode, a cork stopper with a hole in the middle is fixed as an insulated handle. A copper sleeve is strengthened under the cork, to which an electric wire is connected.

The electrical wire used in the appliance must be carefully insulated. Copper couplings and the ends of the wires are wrapped with insulating tape.

When lightly pressing the cork handle, the upper electrode should touch the lower one and, upon cessation of pressure, should return to its original position. The wash bottle with concentrated H 2 SO 4 is connected to a balloon of hydrogen.

The outlet tube passing through the bottom plug of the reaction chamber is connected to a tee. The lower knee of the tee reaches almost to the bottom of the bottle, half filled with water. A short brass tube is attached to the upper knee using a rubber tube with a screw clamp I put on it, into the lower end of which a loose glass wool tampon is inserted. The outlet tube of a bottle of water is connected with a short glass tube using a rubber tube with a clamp II.

The reaction mixture is prepared by grinding in a mortar 6 G calcium orthophosphate, 4 G quartz sand and 3 G coke or charcoal. After calcination over high heat in a closed crucible, the mixture is cooled in a desiccator.

Before the experiment, the mixture is poured into the electrode crucible and pressed against the walls in such a way that an empty space in the form of a cone remains in the middle of the mixture, up to the lower electrode.

Instead of a flask with two tubes, you can use a refractory glass tube with a diameter of about 50 mm. In the absence of a crucible, the reaction mixture can be placed in a conical recess 15 mm made at the upper end of the lower electrode; the carbon electrode in this case should have a diameter of 20 mm. As the top electrode, a carbon electrode with a diameter of 5 mm applied to an electric arc. The experiment is carried out in the dark. Clamp II is closed, clamp I is opened and a strong current of hydrogen is passed through the instrument. After making sure that the hydrogen coming out of the device is pure, they light it at the end of the brass tube and regulate the current so that the flame is calm and not very large. The current is turned on and by pressing on the upper electrode an electric arc is created (10-15 with). After a while, the hydrogen flame turns emerald green (in order to make the color change more noticeable, a porcelain cup is brought into the flame).

Vapors of white phosphorus formed in the reaction vessel are carried away with gases into a flask with water and condense here in the form of small balls. If clamp II is opened and clamp I is closed, then a cold flame of phosphorus can be observed at the end of the vent tube coming out of a bottle of water.

With circular movements of the upper electrode, new portions of the reaction mixture are introduced into the voltaic arc.

To obtain red phosphorus, the hydrogen flow is reduced so that phosphorus vapor does not leave the reaction chamber so quickly.

If you turn off the arc, then on the inner walls of the flask you can notice a coating of red, and on the cold parts of the wall - white phosphorus.

A cold glow or cold flame of phosphorus is observed throughout the entire experiment.

After some cooling of the crucible, the condensation bottle is turned off without stopping the flow of hydrogen.

At the end of the experiment and the complete cooling of the device in a stream of hydrogen, the electrodes are removed, and the flask is left for some time in moist air under draft. To wash the flask, use water with sand or concentrated H 2 SO 4.

Instead of hydrogen, carbon dioxide can be used in the experiment, but the formation of phosphorus in this case is not so effective. The cold glow or cold flame of phosphorus in this case also has a green color.

Small balls of condensed white phosphorus are placed in a flask with cold water and store for future experiments.

Experience. Preparation of white phosphorus by reduction of sodium metaphosphate with aluminum powder in the presence of silicon dioxide. Reaction equation:

6NaPO 3 + 10Al + 3SiO 2 \u003d 6P + 5Al 2 O 3 + 3Na 2 SiO 3.

Recovery is carried out by heating in a refractory tube 25 cm and diameter 1-1.5 cm, connected on one side to a source of pure hydrogen (a cylinder or a Kipp apparatus), and on the other side to a tube through which gaseous products are discharged into a crystallizer with water.

In a refractory tube pour a mixture consisting of 1 wt. including NaRO 3 , 3 wt. including SiO 2 and 0.5 wt. including aluminum filings. With the help of asbestos plugs, the tube is connected on one side through a washing bottle containing concentrated H 2 SO 4 to a source of hydrogen, and on the other - to a drain tube.

After removing air from the device with a strong current of hydrogen and making sure that the outgoing hydrogen is pure, a refractory tube is heated using a Teklu burner with a dovetail. The phosphorus formed by the above reaction is distilled and condensed in the form of small balls in a crystallizer with water. In the dark, you can see the green glow of phosphorus in the tube.

At the end of the experiment, the apparatus is disassembled only after it has been completely cooled in a stream of hydrogen.

The resulting phosphorus is placed for storage in a jar of cold water.

Sodium metaphosphate can be obtained by calcining sodium ammonium hydroorthophosphate hydrate; reaction equation:

NaNH 4 HPO 4 4H 2 O = NaPO 3 + NH 3 + 5H 2 O.

Experience. Obtaining a small amount of white phosphorus from red. The experiment is carried out in a test tube with a length of 17-20 cm and diameter 1.5 cm in an atmosphere of carbon dioxide.

In a test tube, which is held in a vertical position, 0.3-0.5 G dry red phosphorus so that the walls of the test tube remain clean.

The test tube is loosely closed with a rubber stopper with a glass tube reaching almost to the bottom, through which a weak current of carbon dioxide enters the test tube. After filling the tube with carbon dioxide, the glass tube is pulled out so that the tip of the tube remaining in the tube is no longer than 5-6 cm. The test tube at the very hole is fixed in the clamp of the tripod in a horizontal position and the part of it where the phosphorus is located is slightly heated. At the same time, the evaporation of red phosphorus and the precipitation of droplets of white phosphorus on the cold walls of the test tube are observed.

The precipitation of white phosphorus in the dark is clearly visible due to the glow due to slow oxidation. In the dark, the formation of a cold flame (glow) of phosphorus at the opening of the test tube is also observed. If the experiment is carried out in the light, freshly prepared white phosphorus partially turns into red.

At the bottom of the test tube, only impurities contained in phosphorus remain.

At the end of the experiment, the test tube is cooled in a stream of carbon dioxide and tapped on it from time to time to facilitate the solidification of supercooled white phosphorus. After cooling, the test tube with white phosphorus is placed in a beaker of water and heated to 50° to melt all the phosphorus and collect it at the bottom of the tube. After the white phosphorus has solidified, it is extracted by cooling the test tube with a jet cold water. Upon receipt of a very small amount of phosphorus, it is removed from the test tube by burning or heating with a concentrated alkali solution.

To remove traces of phosphorus from the tube through which carbon dioxide was supplied, and the rubber stopper, a solution of KMnO 4 or AgNO 3 is used.

PURIFICATION OF WHITE PHOSPHORUS

White phosphorus can be purified by distillation with water vapor in an atmosphere of carbon dioxide, filtering phosphorus molten in water through suede in an airless space, treatment with a chromium mixture or sodium hypobromite, followed by washing with distilled water.

PHYSICO-CHEMICAL PROPERTIES OF WHITE PHOSPHORUS

Phosphorus is known in several allotropic modifications: white, red, violet and black. In laboratory practice, one has to meet with white and red modifications.

White phosphorus is a solid. AT normal conditions it is yellowish, soft, and wax-like in appearance. It is easily oxidized and flammable. White phosphorus is poisonous - it leaves painful burns on the skin. White phosphorus goes on sale in the form of sticks of different lengths with a diameter of 0.5-2 cm.

White phosphorus is easily oxidized, and therefore it is stored under water in carefully sealed dark glass vessels in poorly lit and not very cold rooms (to avoid cracking the jars due to freezing water). The amount of oxygen contained in water and oxidizing phosphorus is very small; it is 7-14 mg per liter of water.

Under the influence of light, white phosphorus turns into red.

With slow oxidation, the glow of white phosphorus is observed, and with vigorous oxidation, it ignites.

White phosphorus is taken with tweezers or metal tongs; in no case should you touch it with your hands.

In case of a burn with white phosphorus, the burned area is washed with a solution of AgNO 3 (1:1) or KMnO 4 (1:10) and a wet dressing is applied soaked in the same solutions or a 5% solution of copper sulfate, then the wound is washed with water and after smoothing the epidermis, apply vaseline bandage with methyl violet. For severe burns, see a doctor.

Solutions of silver nitrate, potassium permanganate and copper sulfate oxidize white phosphorus and thus stop its damaging effect.

In case of white phosphorus poisoning, take a teaspoon of a 2% copper sulfate solution orally until vomiting occurs. Then, using the Mitcherlich test, based on the luminescence, the presence of phosphorus is determined. For this, water acidified with sulfuric acid is added to the vomiting of the poisoned person, and distilled in the dark; at the content of phosphorus, the glow of vapors is observed. A Wurtz flask is used as a device, to the side tube of which a Liebig condenser is attached, from where the distilled products enter the receiver. If phosphorus vapor is directed into a solution of silver nitrate, then a black precipitate of metallic silver forms, which is formed according to the equation given in the experiment on the reduction of silver salts with white phosphorus.

Already 0.1 G white phosphorus is a lethal dose for an adult.

White phosphorus is cut with a knife or scissors in a porcelain mortar under water. When using water at room temperature, phosphorus crumbles. Therefore, it is better to use warm water, but not higher than 25-30 °. After cutting phosphorus into warm water it is transferred to cold water or cooled with a jet of cold water.

White phosphorus is a highly flammable substance. It ignites at a temperature of 36-60°, depending on the concentration of oxygen in the air. Therefore, when conducting experiments, in order to avoid an accident, it is necessary to take into account every grain of it.

Drying of white phosphorus is carried out by quickly applying thin asbestos or filter paper to it, avoiding friction or pressure.

When phosphorus ignites, it is extinguished with sand, a wet towel or water. If burning phosphorus is on a sheet of paper (or asbestos), this sheet must not be touched, since molten burning phosphorus can be easily spilled.

White phosphorus melts at 44°, boils at 281°. White phosphorus is melted with water, because in contact with air, molten phosphorus ignites. By fusion and subsequent cooling, white phosphorus can be easily recovered from the waste. To do this, white phosphorus waste from various experiments, collected in a porcelain crucible with water, is heated in a water bath. If crust formation is noticeable on the surface of the molten phosphorus, a little HNO 3 or a chromium mixture is added. The crust is oxidized, small grains merge into total weight and after cooling with a jet of cold water one piece of white phosphorus is obtained.

Under no circumstances should phosphorus residues be thrown into the sink, as, accumulating in the bends of the sewer elbow, it can cause burns to maintenance workers.

Experience. Melting and supercooling of molten white phosphorus. A piece of white phosphorus the size of a pea is placed in a test tube with water. The test tube is placed in a beaker filled almost to the top with water and fixed in a vertical position in a tripod clamp. The glass is slightly heated and using a thermometer determine the temperature of the water in the test tube at which phosphorus melts. After the end of melting, the tube is transferred to a beaker with cold water and solidification of phosphorus is observed. If the tube is stationary, then at a temperature below 44° (up to 30°) white phosphorus remains in a liquid state.

The liquid state of white phosphorus, cooled below its melting point, is a state of supercooling.

After the end of the experiment, in order to more easily extract the phosphorus, it is melted again and the test tube is immersed with the hole up in an inclined position in a vessel with cold water.

Experience. Attaching a piece of white phosphorus to the end of the wire. To melt and solidify white phosphorus, a small porcelain crucible with phosphorus and water is used; it is placed in a glass of warm and then cold water. The wire for this purpose is taken iron or copper with a length of 25-30 cm and diameter 0.1-0.3 cm. When the wire is immersed in solidifying phosphorus, it easily attaches to it. In the absence of a crucible, a test tube is used. However, due to the insufficiently even surface of the test tube, it is sometimes necessary to break it in order to extract the phosphorus. To remove white phosphorus from the wire, it is immersed in a glass of warm water.

Experience. Determination of the specific gravity of phosphorus. At 10°, the specific gravity of phosphorus is 1.83. Experience allows us to make sure that white phosphorus is heavier than water and lighter than concentrated H 2 SO 4.

When a small piece of white phosphorus is introduced into a test tube with water and concentrated H 2 SO 4 (specific weight 1.84), it is observed that phosphorus sinks in water, but floats on the surface of the acid, melting due to the heat released when concentrated H 2 SO is dissolved 4 in water.

To pour concentrated H 2 SO 4 into a test tube with water, use a funnel with a long and narrow neck, reaching to the end of the test tube. Pour in the acid and remove the funnel from the test tube carefully so as not to cause mixing of the liquids.

At the end of the experiment, the contents of the test tube are stirred with a glass rod and cooled from the outside with a stream of cold water until the phosphorus solidifies so that it can be removed from the test tube.

When using red phosphorus, it is observed that it sinks not only in water, but also in concentrated H 2 SO 4, since its specific gravity (2.35) is greater than the specific gravity of both water and concentrated sulfuric acid.

WHITE PHOSPHORUS GLOW

Due to the slow oxidation that occurs even at ordinary temperatures, white phosphorus glows in the dark (hence the name "luminiferous"). Around a piece of phosphorus in the dark, a greenish luminous cloud appears, which, when the phosphorus vibrates, is set in a wave-like motion.

Phosphorescence (luminescence of phosphorus) is explained by the slow oxidation of phosphorus vapor by oxygen in the air to phosphorous and phosphorus anhydride with the release of light, but without the release of heat. In this case, ozone is released, and the air around is ionized (see the experiment showing the slow combustion of white phosphorus).

Phosphorescence depends on temperature and oxygen concentration. At 10°C and normal pressure, phosphorescence proceeds weakly, and in the absence of air it does not occur at all.

Substances that react with ozone (H 2 S, SO 2, Cl 2, NH 3, C 2 H 4, turpentine oil) weaken or completely stop phosphorescence.

The conversion of chemical energy into light energy is called "chemiluminescence".

Experience. Observation of the glow of white phosphorus. If you observe in the dark a piece of white phosphorus in a glass and not completely covered with water, then a greenish glow is noticeable. In this case, wet phosphorus slowly oxidizes, but does not ignite, since the temperature of the water is below the flash point of white phosphorus.

The glow of white phosphorus can be observed after a piece of white phosphorus has been exposed to air for a short time. If you put a few pieces of white phosphorus in a flask on glass wool and fill the flask with carbon dioxide, lowering the end of the outlet tube to the bottom of the flask under glass wool, and then slightly heat the flask by dipping it into a vessel with warm water, then in the dark you can observe the formation of a cold pale greenish flame (you can safely put your hand into it).

The formation of a cold flame is explained by the fact that carbon dioxide leaving the flask entrains phosphorus vapor, which begins to oxidize when it comes into contact with air at the opening of the flask. In a flask, white phosphorus does not ignite, because it is in an atmosphere of carbon dioxide. At the end of the experiment, the flask is filled with water.

In describing the experiment for the production of white phosphorus in an atmosphere of hydrogen or carbon dioxide, it was already mentioned that carrying out these experiments in the dark makes it possible to observe the glow of white phosphorus.

If you make an inscription with phosphorus chalk on a wall, a sheet of cardboard or paper, then thanks to phosphorescence, the inscription remains visible for a long time in the dark.

Such an inscription cannot be made on a blackboard, because after that ordinary chalk does not stick to it and the board has to be washed with gasoline or another stearin solvent.

Phosphorus chalk is obtained by dissolving liquid white phosphorus in molten stearin or paraffin. To do this, approximately two parts by weight of stearin (candle pieces) or paraffin are added to a test tube to one weight part of dry white phosphorus, the test tube is covered with cotton wool to prevent oxygen from entering, and heated with continuous shaking. After the end of melting, the test tube is cooled with a jet of cold water, then the test tube is broken and the solidified mass is removed.

Phosphorus chalk is stored under water. When using a piece of such chalk is wrapped in wet paper.

Phosphorus chalk can also be obtained by adding small pieces of dried white phosphorus to paraffin (stearin) melted in a porcelain cup. If the paraffin ignites when phosphorus is added, it is extinguished by covering the cup with a piece of cardboard or asbestos.

After some cooling, the solution of phosphorus in paraffin is poured into dry and clean test tubes and cooled with a stream of cold water until it solidifies into a solid mass.

After that, test tubes are broken, chalk is removed and stored under water.

SOLUBILITY OF WHITE PHOSPHORUS

In water, white phosphorus is sparingly soluble, slightly soluble in alcohol, ether, benzene, xylene, methyl iodide and glycerin; dissolves well in carbon disulfide, sulfur chloride, phosphorus trichloride and tribromide, carbon tetrachloride.

Experience. Dissolution of white phosphorus in carbon disulfide. Carbon disulfide is a colorless, highly volatile, highly flammable, poisonous liquid. Therefore, when working with it, avoid inhaling its vapors and turn off all gas burners.

Three or four pieces of white phosphorus the size of a pea are dissolved with light shaking in a glass of 10-15 ml carbon disulfide.

If a small sheet of filter paper is moistened with this solution and held in air, the paper ignites after a while. This is because carbon disulfide evaporates quickly, and the finely divided white phosphorus remaining on the paper quickly oxidizes at ordinary temperatures and ignites due to the heat released during oxidation. (It is known that the ignition temperature of various substances depends on the degree of their grinding.) It happens that paper does not ignite, but only chars. Paper moistened with a solution of phosphorus in carbon disulfide is kept in air with metal tongs.

The experiment is carried out carefully so that drops of a solution of phosphorus in carbon disulfide do not fall on the floor, on the table, on clothes or on hands.

If the solution gets on the hand, it is quickly washed with soap and water, and then with a solution of KMnO 4 (to oxidize particles of white phosphorus that have fallen on the hands).

The solution of phosphorus in carbon disulfide remaining after the experiments is not stored in the laboratory, since it can easily ignite.

WHITE PHOSPHORUS TRANSFORMATION TO RED

White phosphorus is converted to red according to the equation:

P (white) = P (red) + 4 kcal.

The process of converting white phosphorus to red is greatly accelerated by heating, under the influence of light and in the presence of traces of iodine (1 G iodine at 400 G white phosphorus). Iodine, combining with phosphorus, forms phosphorus iodide, in which white phosphorus dissolves and quickly turns into red with the release of heat.

Red phosphorus is obtained by prolonged heating of white phosphorus in a closed vessel in the presence of traces of iodine to 280-340 °

With long-term storage of white phosphorus in the light, it gradually turns into red.

Experience. Obtaining a small amount of red phosphorus from white. In a glass tube 10-12 long, closed at one end cm and diameter 0.6-0.8 cm they introduce a piece of white phosphorus the size of a grain of wheat and a very small crystal of iodine. The tube is sealed and suspended in an air bath over a tray of sand, then heated to 280-340° and the transformation of white phosphorus into red is observed.

Partial conversion of white phosphorus to red can also be observed by slightly heating a test tube with a small piece of white phosphorus and a very small crystal of iodine. Before starting heating, the test tube is closed with a swab of glass (asbestos or ordinary) wool and a tray with sand is placed under the test tube. The tube is heated for 10-15 minutes (without bringing the phosphorus to a boil) and the transformation of white phosphorus into red is observed.

White phosphorus remaining in the test tube can be removed by heating with a concentrated alkali solution or by burning.

The transformation of white phosphorus into red can also be observed by heating a small piece of phosphorus in a test tube in an atmosphere of carbon dioxide to a temperature below boiling.

COMBUSTION OF WHITE PHOSPHORUS

When white phosphorus burns, phosphoric anhydride is formed:

P 4 + 5O 2 \u003d 2P 2 O 5 + 2 x 358.4 kcal.

You can observe the combustion of phosphorus in air (slow and fast) and under water.

Experience. Slow combustion of white phosphorus and air composition. This experiment was not described as a way to obtain nitrogen, since it does not completely bind the oxygen contained in the air.

The slow oxidation of white phosphorus by atmospheric oxygen occurs in two stages; in the first stage, phosphorous anhydride and ozone are formed according to the equations:

2P + 2O 2 \u003d P 2 O 3 + O, O + O 2 \u003d O 3.

In the second stage, phosphorous anhydride is oxidized to phosphoric anhydride.

The slow oxidation of white phosphorus is accompanied by luminescence and ionization of the surrounding air.

An experiment showing the slow burning of white phosphorus should last at least three hours. The apparatus required for the experiment is shown in Fig.

In a cylinder expanded at the opening, almost filled with water, a graduated tube with a closed end, containing about 10 ml water. Tube length 70 cm, diameter 1.5-2 cm. After lowering the graduated tube, remove the finger from the tube opening, bring the water in the tube and cylinder to the same level, and note the volume of air contained in the tube. Without raising the tube above the water level in the cylinder (in order not to let in additional air), a piece of white phosphorus fixed at the end of the wire is introduced into the air space of the tube.

After three to four hours, or even after two or three days, a rise in water in the tube is noted.

At the end of the experiment, the wire with phosphorus is removed from the tube (without raising the tube above the water level in the cylinder), the water in the tube and cylinder is brought to the same level, and the volume of air remaining after the slow oxidation of white phosphorus is noted.

Experience shows that as a result of the binding of oxygen by phosphorus, the volume of air has decreased by one fifth, which corresponds to the oxygen content in the air.

Experience. Rapid combustion of white phosphorus. In view of the fact that during the reaction of the combination of phosphorus with oxygen, a large number of heat, in air, white phosphorus ignites spontaneously and burns with a bright yellowish-white flame, forming phosphoric anhydride, a white solid that combines very vigorously with water.

It was already mentioned earlier that white phosphorus ignites at 36-60 °. To observe its self-ignition and combustion, a piece of white phosphorus is placed on a sheet of asbestos and covered with a glass bell or a large funnel, on the neck of which a test tube is put on.

Phosphorus can be easily set on fire with a glass rod heated in hot water.

Experience. Comparison of ignition temperatures of white and red phosphorus. At one end of a copper plate (length 25 cm, width 2.5 cm and thickness 1 mm) put a small piece of dried white phosphorus, pour a small pile of red phosphorus on the other end. The plate is placed on a tripod and at the same time approximately equally burning gas burners are brought to both ends of the plate.

White phosphorus ignites immediately, and red phosphorus only when its temperature reaches approximately 240°.

Experience. Ignition of white phosphorus under water. A test tube with water containing several small pieces of white phosphorus is dipped into a beaker containing hot water. When the water in the test tube is heated to 30-50°C, a current of oxygen is passed into it through the tube. Phosphorus ignites and burns, scattering bright sparks.

If the experiment is carried out in the beaker itself (without a test tube), the beaker is placed on a tripod mounted on a tray of sand.

REDUCTION OF SILVER AND COPPER SALTS WITH WHITE PHOSPHORUS

Experience. When a piece of white phosphorus is introduced into a test tube with a solution of silver nitrate, a precipitate of metallic silver is observed (white phosphorus is an energetic reducing agent):

P + 5AgNO 3 + 4H 2 O \u003d H 3 RO 4 + 5Ag + 5HNO 3.

If white phosphorus is introduced into a test tube with a solution of copper sulfate, then metallic copper precipitates:

2P + 5CuSO 4 + 8H 2 O \u003d 2H 3 PO 4 + 5H 2 SO 4 + 5Cu.

RED PHOSPHORUS

Methods for obtaining red phosphorus from white are described above.

IMPURITIES

Red phosphorus contains traces of white phosphorus, phosphoric and pyrophosphoric acids.

The presence of phosphoric acid is explained by the combination of phosphoric anhydride with air moisture, and the formation of phosphoric anhydride is explained by the slow oxidation of traces of white phosphorus. When wet phosphorus is oxidized with oxygen, in addition to phosphorous and phosphoric anhydrides, hypophosphorous acid is also formed.

CLEANING AND STORING RED PHOSPHORUS

Red phosphorus is purified by boiling with a dilute NaOH solution, after which it is thoroughly washed by decantation, and then on a filter with distilled water.

The washed phosphorus is dried with filter paper, placed on a watch glass and kept in an oven at 105°.

Store it in jars closed with paraffin cork.

PROPERTIES

Red phosphorus is a powder (sp. weight 2.35), insoluble in water and carbon disulfide, sublimes at 416° and ignites at 240°. Unlike white, red phosphorus is not poisonous.

The sublimation temperature of red phosphorus is determined in an atmosphere of carbon dioxide. Vapors of red phosphorus, thickening, give white phosphorus.

Red phosphorus is chemically less active than white phosphorus. It does not glow in air and in oxygen, but glows in an ozone atmosphere; does not displace metals (copper, silver, etc.) from their salts; indifferent to alkalis; reacts with halogens, oxygen and sulfur at a higher temperature than white phosphorus.

Experience. An explosion of a mixture of red phosphorus and bartholium salt. When picking up red phosphorus powder, you need to be careful, as it can ignite from friction.

To conduct the experiment, a small amount of a mixture of red phosphorus and bartholite salt is poured onto an anvil, a piece of rail or a stone and hit with a hammer.

In order to avoid injury, in no case should you take a large amount of the mixture.

The powders are mixed gently by simply rocking the sheet. For one part of dry powder of red phosphorus, take at least two parts of berthollet salt powder. During the experiment, turn Special attention on the composition of the mixture, its quantity, so that the explosion is not very strong, and also so that the mixture does not explode unexpectedly in the hands of the experimenter.

An excess of red phosphorus leads to the fact that during the experiment, phosphorus simply ignites; with wet phosphorus, the experiment fails.

Experience. An explosion of a mixture of red phosphorus, bartholium salt and sulfur. On a piece of paper carefully mix 0.2-0.3 G dry powder of red phosphorus, 2-3 G dry powder of Berthollet salt and 0.5 G sulfur powder.

When mixing, a piece of paper is held with both hands, alternately moving them up and down a little. received homogeneous mixture divided into 5-6 parts.

One part of the mixture is poured onto a piece of paper 10x10 cm, put a pellet in it, fold the corners of the paper and twist them lightly together.

The resulting knot is thrown onto something solid (stone or cement floor) - a strong explosion occurs.

If at least one of the starting materials was wet, the experiment fails.

APPLICATIONS OF PHOSPHORUS

White phosphorus is used for the production of hydrogen phosphide, phosphides, phosphoric acid, some pharmaceuticals, aniline dyes, smoke-forming and incendiary liquids, for the formation of smoke screens, and as a poison against rats.

Previously, white phosphorus was used in match production; at present it is not used for this purpose, because it is poisonous and flammable.

Currently, match production uses red phosphorus. For a match head, a mixture is prepared next composition(in wt%):

Bertoletova salt 46.5
Minium or mummy 15.3
Chrome peak 1.5
Ground glass 17.2
Sulfur 4.2
Bone glue 11.5
Zinc white 3.8

The matchbox spread contains 30.8 wt. % red phosphorus.

For better ignition of the match, it is impregnated with paraffin, and so that after extinguishing it does not smolder - with sodium phosphate.

Red phosphorus is used for the production of hydrogen bromide and iodide, phosphorus compounds with halogens, organic dyes, for the production of phosphorous bronzes (having a high viscosity) and for filling incendiary shells.

PHOSPHORUS COMPOUNDS

PHOSPHORUS HYDROGEN PH 3 (PHOSPHINE)

SPREAD

Phosphorous hydrogen is formed during the decomposition of organic substances containing phosphorus.

RECEIVING

Phosphoric hydrogen is a very poisonous gas, so all experiments with it are carried out under traction.

Experience. Obtaining hydrogen phosphide by heating white phosphorus with a 30-50% KOH solution. Reaction equation:

4P + 3KOH + 3H 2 O \u003d PH 3 + 3KN 2 RO 2.

With this production method, in addition to gaseous hydrogen phosphide, liquid hydrogen phosphide, gaseous hydrogen and potassium acid hypophosphite are also formed according to the equations:

6P + 4KOH + 4H 2 O \u003d P 2 H 4 + 4KN 2 PO 2,

2P + 2KOH + 2H 2 O \u003d H 2 + 2KN 2 PO 2.

Liquid hydrogen phosphide, interacting with potassium hydroxide in an aqueous medium, forms gaseous hydrogen phosphide, hydrogen and potassium acid hypophosphite according to the equations:

2P 2 H 4 + KOH + H 2 O \u003d ZRN 3 + KN 2 RO 2,

R 2 H 4 + 2KOH + 2H 2 O \u003d ZN 2 + 2KN 2 RO 2.

Acid potassium hypophosphite in alkaline environment turns into potassium orthophosphate with the release of hydrogen:

KN 2 PO 2 + 2KOH \u003d 2H 2 + K 3 PO 4.

According to the above reaction equations, when white phosphorus is heated with potassium hydroxide, gaseous hydrogen phosphide, hydrogen and potassium orthophosphate are formed.

The hydrogen phosphorous obtained in this way spontaneously ignites. This is because it contains some vapors of self-igniting liquid hydrogen phosphide and hydrogen.

Instead of potassium oxide hydrate, sodium, calcium or barium oxide hydrates can be used. Reactions with them proceed similarly.

The device is a round-bottom flask with a capacity of 100-250 ml, tightly closed with a rubber stopper, through which a tube must be passed, directing gaseous products into the crystallizer with water.

The flask is filled to 3/4 of its volume with a 30-50% KOH solution, into which 2-3 pea-sized pieces of white phosphorus are thrown. The flask is fixed in a tripod clamp and connected to a crystallizer filled with water using a drain tube (Fig.).

When the flask is heated, potassium hydroxide reacts with white phosphorus according to the above reaction equations.

Liquid hydrogen phosphide, having reached the surface of the liquid in the flask, immediately ignites and burns in the form of sparks; this continues until the remaining oxygen in the flask is used up.

When the flask is strongly heated, liquid hydrogen phosphide is distilled and ignites gaseous hydrogen phosphide and hydrogen over water. Phosphoric hydrogen burns with a yellow flame, forming phosphorus anhydride in the form of white smoke rings.

At the end of the experiment, reduce the flame under the flask, remove the plug with the outlet tube, stop heating and leave the device under draft until it is completely cooled.

Unused phosphorus is thoroughly washed with water and stored for the next experiments.

Experience. Preparation of (spontaneously flammable) gaseous hydrogen phosphide by decomposition of calcium phosphide with water. The reaction proceeds according to the equation:

Ca 3 P 2 + 6H 2 O \u003d 2PH 3 + 3Ca (OH) 2.

The following reactions also take place simultaneously:

Ca 3 P 2 + 6H 2 O \u003d P 2 H 4 + H 2 + 3Ca (OH) 2,

4P 2 H 4 + Ca (OH) 2 + 2H 2 O \u003d 6PH 3 + Ca (H 2 PO 2) 2,

P 2 H 4 + Ca (OH) 2 + 2H 2 O \u003d 3H 2 + Ca (H 2 RO 2) 2.

The device is a small flask with a straight outlet tube and a large beaker.

For weighting in a flask with a capacity of 100 ml pour lead shot, then add a small amount of dry calcium phosphide and a few drops of ether. The flask is closed with a rubber stopper, through which a straight glass tube 7-8 cm and diameter 3-5 mm starting at the bottom edge of the cork. Having put several lead rings on the neck of the flask, a rope is tied to it. After holding the flask for some time in the palm of your hand to evaporate the ether, it is immersed on a string in a large glass (with a capacity of about 3 l) with water. First, air bubbles and ether vapors are released from the flask, then, when the gas pressure in the flask decreases, a small amount of water enters the flask and the decomposition of calcium phosphide begins.

The gaseous products formed as a result of the decomposition of calcium phosphide prevent the continuous flow of water into the flask.

As the resulting gases reach the surface of the water, they flare up and, burning, form phosphoric anhydride in the form of white smoke rings.

Water enters the flask in small portions at the moment of decreasing gas pressure and forms hydrogen phosphide until calcium phosphide is completely consumed.

Lead shot and rings are used to immerse the flask in a glass of water.

This experiment can be carried out in another way. A few pieces of calcium phosphide are thrown into a glass of water. The gas bubbles released during the decomposition of calcium phosphide ignite when leaving the water. When hydrogen phosphorous is burned, phosphoric anhydride is formed, which in this case also rises above the glass in the form of rings of white smoke.

Calcium phosphide is taken with tweezers or tongs.

Obtaining pure (spontaneously non-flammable) hydrogen phosphide is described in the section on the properties of diphosphine.

Experience. Preparation of hydrogen phosphide by the action of dilute HCl and H 2 SO 4 (or water acidified with one of these acids) on calcium, zinc, magnesium and aluminum phosphides. Reaction equations:

Me 3 P 2 + 6HCl \u003d 2PH 3 + 3MeCl 2,

Me - Ca, Mg, Zn,

AlP + 3HCl = PH 3 + AlCl 3.

In this experiment, along with gaseous phosphorous hydrogen, liquid phosphorous hydrogen and gaseous hydrogen are formed.

One of the phosphides listed above is added to a beaker with dilute HCl (sp. weight 1.12) or dilute H 2 SO 4 . The evolution of hydrogen phosphide is observed, spontaneously igniting over the solution in the beaker.

Experience. Obtaining pure phosphorous hydrogen PH 3 by decomposition of phosphorous and hypophosphorous acids. When heated, the following reactions take place:

4H 3 RO 3 \u003d PH 3 + 3H 3 RO 4,

2H 3 RO 2 \u003d PH 3 + H 3 RO 4.

Concentrated acid solutions are heated in small glass flasks. The evolved gaseous products are sent through a tube to a crystallizer with water.

Experience. Preparation of pure gaseous hydrogen phosphide by the action of a dilute solution of potassium hydroxide on phosphonium iodide. Reaction equation:

PH 4 I + KOH \u003d PH 3 + KI + H 2 O.

To obtain hydrogen phosphide, a solution of KOH is added from a dropping funnel to a Wurtz flask with small glass tubes and dry pH 4 I.

PRODUCTION AND PROPERTIES OF PHOSPHONIUM IODIDE

Dissolve in carbon disulfide 50 G white phosphorus. Gradually add 65 G iodine. After removal of carbon disulfide by evaporation, crystals of phosphorus iodide P 2 I 4 remain; they are placed in a Wurtz flask with a wide side tube. A weak current of CO 2 is passed through the Wurtz flask, and then water is poured from the dropping funnel.

As a result, phosphorous acid, a small amount of free hydrogen iodide and phosphonium iodide are formed in the Wurtz flask. When heated to 80°, the latter sublimates and can be collected in a wide tube cooled from the outside. The resulting phosphonium iodide is a colorless crystalline substance that decomposes with water.

We have already met with the formation of phosphonium iodide in experiments on the production of hydrogen iodide.

PROPERTIES OF GASEOUS PHOSPHORUS HYDROGEN

Under normal conditions, gaseous hydrogen phosphide is a colorless, highly toxic gas with bad smell rotten fish (or garlic). It is highly soluble in water (under normal conditions in 5 l water dissolves 1 l pH 3), but does not chemically interact with it. It is poorly soluble in alcohol and ether. When cooled, it thickens into a liquid, which boils at -87.4° and solidifies into a crystalline mass at -132.5°. Critical temperature of hydrogen phosphide 52.8°, critical pressure 64 atm.

Phosphoric hydrogen is a very strong reducing agent; ignites in air at 150° and burns with a yellow flame to form phosphoric anhydride according to the equation:

2РН 3 + 4O 2 = Р 2 O 5 + 3Н 2 O

The combustion of gaseous hydrogen phosphide has already been discussed in experiments on its production.

Experience. Recovery of aqueous solutions of silver and copper salts with gaseous hydrogen phosphorous. Reaction equations:

6AgNO 3 + PH 3 + 3H 2 O \u003d 6HNO 3 + H 3 PO 3 + 6Ag,

3CuSO 4 + PH 3 + 3H 2 O \u003d 3H 2 SO 4 + H 3 PO 3 + 3Cu.

The experiment is carried out in test tubes. As a result of the reaction, not only silver and copper are released, but the corresponding phosphides are also formed, for example:

3СuSO 4 + 2РН 3 = Сu 3 Р 2 + 3Н 2 SO 4

Copper salts (CuSO 4 and Cu 2 Cl 2) absorb gaseous hydrogen phosphide, and this is used to separate the gaseous mixture of hydrogen phosphide and hydrogen - it is passed through washing vessels with copper salts.

Gaseous hydrogen phosphorous also reduces nitric, sulfuric and sulphurous acids, gold salts and other compounds.

The interaction of gaseous hydrogen phosphide with chlorine has already been discussed in the description of experiments to study the properties of chlorine.

Gaseous hydrogen phosphide combines directly with hydrohalic acids to form phosphonium salts (obtaining phosphonium iodide is described above). Equal volumes of hydrogen iodide and hydrogen phosphide combine to form colorless cubic crystals of phosphonium iodide.

CALCIUM PHOSPIDE

Experience. Preparation and properties of calcium phosphide. Calcium phosphide is obtained from small chips of calcium and red phosphorus under draft. White phosphorus is not used for this purpose, since the reaction with it proceeds too violently.

The device is a glass tube with a length of 10-12 cm and diameter 0.5 cm fixed at one end in the tripod clamp horizontally. Mixture 1 is placed in the middle of the tube G small chips of calcium and 1 G dry red phosphorus. When the tube is heated, a violent combination of both substances occurs with the formation of Ca 3 P 2 - a light brown solid. After cooling, the tube is broken with a pestle in a large mortar. Calcium phosphide is taken from the mortar with a spatula, tweezers or metal tongs and placed in a dry jar for storage. The jar is tightly closed and filled with paraffin to prevent the decomposition of calcium phosphide under the influence of atmospheric moisture.

All fragments of the tube contaminated with calcium phosphide are also carefully removed, since toxic products are formed during the decomposition of the latter.

The interaction of calcium phosphide with water and dilute acids was considered in experiments on the production of gaseous hydrogen phosphide.

LIQUID PHOSPHORUS HYDROGEN R 2 H 4 (DIPHOSPHINE)

Usually, diphosphine is formed as a by-product during the production of phosphine, in particular, this occurs when phosphides are decomposed by water. But due to the large difference between the boiling and melting points of phosphine and diphosphine, they can be easily separated by passing the gas mixture through a tube cooled to 0°.

Obtaining diphosphine is carried out in a dark room, since it decomposes under the action of light.

Experience. Preparation and properties of diphosphine. The device is assembled in accordance with fig. A three-necked flask is connected on one side to a long outlet tube passing through a cooling mixture of ice and table salt, and on the other side to a safety tube, the end of which must be lowered into a vessel with water. A three-necked flask is filled to 2/8 of its volume with water and placed in a water bath, with the help of which the temperature of the water in the flask is maintained at a level of about 50 °. A wide straight tube is inserted into the middle neck of a three-necked flask, the upper end of which is closed with a rubber stopper.

Before the start of the experiment, the safety tube is connected to a source of CO 2 to force air out of the instrument. This is done in order to prevent an explosion that can occur during the experiment if there is air in the flask.

After removing air from the device, the free end of the outlet tube is closed with a rubber stopper, the source of CO 2 is disconnected, and the end of the safety tube is lowered into a vessel with water.

A few pieces of calcium phosphide are introduced into the flask through the middle tube and the tube is closed with a rubber stopper.

Phosphoric hydrogen, formed during the decomposition of calcium phosphide, displaces carbon dioxide from the bottle through the safety tube.

After removing carbon dioxide from the flask, remove the cork from the outlet tube. Now the vapors of liquid hydrogen phosphide with the water vapor entrained by them rush into the outlet tube and condense in that part of it that is immersed in the cooling mixture. When this part of the tube is clogged with condensed vapors of hydrogen phosphide and water, the gases again rush into the safety tube.

The free end of the outlet tube with frozen diphosphine is sealed with a gas burner, then the tube is disconnected from the device and the other end is sealed.

Diphosphine under normal conditions is a colorless, water-immiscible liquid, boiling at 51.7° and solidifying at -99°. This liquid ignites spontaneously and burns with a very bright flame, therefore it is stored in the absence of air.

Diphosphine strongly refracts light and does not wet glass walls.

Under the influence of atomized solids, turpentine, heat (30°), light and concentrated HCl, diphosphine decomposes into phosphine and phosphorus according to the equation:

3P 2 H 4 \u003d 4RN 3 + 2P.

Phosphorus absorbs some of the phosphine, forming a compound called solid hydrogen phosphorous.

Taking advantage of the fact that diphosphine decomposes in the presence of concentrated HCl, it is possible to obtain gaseous spontaneously non-flammable hydrogen phosphide. To do this, a mixture of gaseous hydrogen phosphide with vapors of liquid hydrogen phosphide is passed through a wash bottle with concentrated HCl. In this case, solid hydrogen phosphorous remains in the washing flask - a light yellow substance that decomposes under the influence of light into hydrogen and red phosphorus.

Experience. Obtaining pure, spontaneously non-flammable phosphorous hydrogen. The device is assembled according to Fig. The first three-necked flask is filled 2/3 with dilute HCl, the second is filled with concentrated HCl, and water is poured into the crystallizer. The device is assembled and air is removed from it with the help of carbon dioxide, which enters the first three-necked flask. After removing the air, close clamp I on the rubber tube.

After adding calcium phosphide through the middle tube into the first three-necked flask, a mixture of phosphine and diphosphine is formed.

Passing through concentrated HCl, diphosphine decomposes, and pure gaseous hydrogen phosphorous enters the crystallizer with water, which is collected in various vessels according to the water displacement method.

OXYGEN COMPOUNDS OF PHOSPHORUS

Experience. Obtaining and properties of phosphorous anhydride (phosphorus trioxide). Phosphoric anhydride is obtained by passing dry air through heated red phosphorus. Three glass tubes ground to each other serve as a device. The first tube, fixed horizontally in the tripod clamp, serves to heat the red phosphorus. In the second tube, also fixed in a horizontal position, heated to approximately 50 °, a glass wool swab is placed to trap the incoming phosphorus and phosphorus anhydride from the first tube. The third tube is curved, its end is lowered almost to the bottom of a small flask cooled from the outside, in which phosphorus anhydride is condensed.

Phosphoric anhydride - white, crystalline, very wax-like poisonous substance, melting at 23.8° and boiling at 173.1°. (The boiling point can be set by heating phosphorous anhydride under nitrogen.)

Phosphoric anhydride has reducing properties. Heated to 70 °, it ignites and burns out, turning into phosphoric anhydride according to the equation:

P 2 O 3 + O 2 \u003d P 2 O 5.

Gradually, this oxidation, accompanied by luminescence, begins to proceed even at ordinary temperatures.

Phosphoric anhydride forms dimerized P 4 O 10 molecules.

When heated above 210 ° or under the influence of light, phosphorous anhydride decomposes:

2P 4 O 6 \u003d 2P + 3P 2 O 4.

Phosphorous anhydride combines with cold water very slowly, forming phosphorous acid H 3 PO 3. It reacts violently with hot water, forming phosphine and phosphoric acid according to the equation:

P 4 O 6 + 6H 2 O \u003d PH 3 + 3H 3 PO 4.

Experience. Preparation and properties of phosphoric anhydride P 2 O 5 (phosphorus pentoxide). To obtain phosphoric anhydride by burning phosphorus, use the device shown in Fig.

A wide straight glass tube is inserted into the neck of the flask on a rubber stopper, to the end of which a small porcelain crucible is tied with a wire. The tube serves to introduce phosphorus into the crucible and ignite it with a heated wire. Through one of the side tubes, air enters the flask, which, for cleaning, first passes through washing flasks with concentrated solutions of NaOH and H 2 SO 4 . Oxygen-deprived air escapes from the flask through the second tube, carrying phosphoric anhydride with it, condensing in a dry and cold flask. The latter is connected to a water jet pump through a wash bottle with water.

To carry out the experiment, a water-jet pump is turned on, pieces of phosphorus are introduced into the crucible and set on fire. After igniting the phosphorus, the heated wire is removed and the upper end of the wide glass tube is closed with a rubber stopper.

All tubes and plugs in the device must be tightly connected.

Phosphorus burns according to the equation:

4P + 5O 2 \u003d 2P 2 O 5 + 2 x 358.4 kcal.

The resulting phosphoric anhydride condenses in a cold bottle in the form of flakes resembling snow.

The preparation of phosphoric anhydride has already been discussed in the study of the properties of oxygen and phosphorus.

Phosphoric anhydride is purified from impurities of lower oxides of phosphorus by sublimation in a stream of oxygen in the presence of spongy platinum. Store phosphoric anhydride in dry, tightly closed and paraffin-filled jars.

Phosphoric anhydride has the appearance of a white crystalline snow-like substance, but may be amorphous and glassy.

Depending on the number of water molecules attached to the phosphoric anhydride molecule, meta-, pyro- and orthophosphoric acids are formed:

P 2 O 5 + H 2 O \u003d 2HPO 3,

P 2 O 5 + 2H 2 O \u003d H 4 P 2 O 7,

P 2 O 5 + 3H 2 O \u003d 2H 3 PO 4.

Phosphoric anhydride is the most powerful dehydrating agent for gases, so it is filled with drying columns and towers, applying it to asbestos or glass wool. In some cases, it can take away elements of water from other compounds, so it is used in the production of nitric, sulfuric anhydride and other compounds. In air, phosphoric anhydride, attracting moisture, quickly spreads (it should be stored in the absence of moisture).

When phosphoric anhydride comes into contact with water, a violent hydration reaction occurs, accompanied by a strong whistling noise. Dream large quantity In cold water, it gives metaphosphoric acid, and with a large amount of warm water, it forms orthophosphoric acid.

Phosphoric anhydride heated to 250° sublimes and settles on the cold walls of the vessel in the form of monoclinic crystals. When heated in a closed device to 440°, it polymerizes and passes into a powder form, and at 600° it acquires a vitreous form. As a result of vapor condensation, a crystalline form is formed. Phosphoric anhydride melts at 563°.

Experience. Obtaining and properties of metaphosphoric acid HPO 3. In a small glass containing 50 ml of water, add 1-2 tablespoons of phosphoric anhydride. Water becomes cloudy due to the formation of metaphosphoric acid. The solution becomes light if allowed to stand, shake or slightly warm.

When the solution is evaporated, metaphosphoric acid is released in the form of a transparent, ice-like, colorless glassy mass.

Store metaphosphoric acid in jars closed with a paraffin stopper; in the presence of air, it becomes covered with a white coating, which can be removed by washing.

Monobasic metaphosphoric acid refers to acids of medium strength. It is soluble in water. With an excess of water, it passes into pyro- and orthophosphoric acids.

Metaphosphoric acid or mstaphosphate solution with the addition of acetic acid coagulate albumin. You can conduct an experiment in a test tube showing the coagulation of egg white.

Experience. Obtaining and properties of orthophosphoric acid. On the preparation of pure orthophosphoric acid by the oxidation of phosphorus nitric acid mentioned in the study of the properties of nitric acid.

Orthophosphoric acid can also be obtained by heating or long-term storage of metaphosphoric acid, heating phosphorous acid, the action of water on phosphorus pentachloride, phosphorus oxychloride or phosphoric anhydride, and the action of concentrated sulfuric acid on calcium orthophosphate.

Orthophosphoric acid is formed by the action of sulfuric acid on bone ash:

Ca 3 (PO 4) 2 + 3H 2 SO 4 \u003d 3CaSO 4 + 2H 3 PO 4.

In a porcelain cup for 4-5 minutes, heat 5 G bone ash, 5 ml water and 5 ml concentrated H 2 SO 4 (sp. weight 1.84). The contents of the cup are then transferred to a beaker and, after cooling, diluted with an equal volume of cold water.

After filtering the calcium sulfate precipitate and evaporating the clear solution (by heating to 150°C), it thickens, acquiring the consistency of a thick syrup.

If part of the filtered solution is neutralized in the presence of litmus with ammonia (adding it in a small excess), and then silver nitrate is added, a yellow precipitate of silver orthophosphate Ag 3 PO 4 precipitates.

Orthophosphoric acid is a colorless, transparent and solid rhombic crystals, deliquescent in air. It is a tribasic acid of medium strength. It dissolves very easily in water with the release of a small amount of heat. It goes on sale in the form of a 40-95% aqueous solution.

As a result of the replacement of one, two or three hydrogen ions with metals, phosphoric acid forms three series of salts (NaH 2 PO 4 - primary sodium phosphate, Na 2 HPO 4 - secondary - sodium phosphate and Na 3 PO 4 - tertiary sodium phosphate).

The weaker but less volatile phosphoric acid can displace nitric and sulfuric acids from their compounds.

When orthophosphoric acid is heated to 215°, pyrophosphoric acid is obtained in the form of a vitreous mass. The reaction proceeds according to the equation:

2H 3 RO 4 + 35 kcal\u003d H 4 P 2 O 7 + H 2 O,

and when heated above 300 °, pyrophosphoric acid turns into metaphosphoric:

H 4 P 2 O 7 + 6 kcal\u003d 2HPO 3 + H 2 O.

Experience. Preparation and properties of phosphorous acid. The preparation of phosphorous acid by the hydrolysis of phosphorus tribromide, triiodide, and trichloride has been described in experiments on the production of hydrogen bromide and hydrogen iodide and will be touched upon further in experiments on the properties of phosphorus trichloride.

Phosphorous acid is a dibasic acid of medium strength; it forms two series of salts, for example NaH 2 PO 3 - acidic sodium phosphite and Na 2 HPO 3 - average sodium phosphite.

In the free state, H 3 PO 3 is a colorless crystal, spreading in air and easily soluble in water.

When heated, phosphorous acid decomposes into orthophosphoric acid and phosphine according to the equation:

4H 3 RO 3 \u003d 3H 3 RO 4 + PH 3.

Phosphorous acid is a strong reducing agent; when heated, it reduces the solution of mercuric chloride to chloride and even to metallic mercury, and metallic silver is isolated from a solution of silver nitrate:

H 3 RO 3 + 2HgCl 2 + H 2 O \u003d Hg 2 Cl 2 + H 3 RO 4 + 2HCl,

H 3 PO 3 + HgCl 2 + H 2 O \u003d Hg + H 3 RO 4 + HCl,

H 3 PO 3 + 2AgNO 3 + H 2 O \u003d 2Ag + H 3 PO 4 + 2HNO 3.

Experience. The reducing nature of hypophosphorous acid H 3 PO 2. Phosphorous acid and its salts (hypophosphites) reduce salts of copper, silver, mercury, gold and bismuth to the corresponding metals. For example, if a solution of hypophosphorous acid is added to a solution of copper sulfate or silver nitrate, metallic copper, metallic silver is released and orthophosphoric acid is formed according to the equations:

H 3 PO 2 + 2CuSO 4 + 2H 2 O \u003d 2Cu + H 3 PO 4 + 2H 2 SO 4,

H 3 PO 2 + 4AgNO 3 + 2H 2 O \u003d 4Ag + H 3 PO 4 + 4HNO 3.

Phosphorous acid reduces bromine and iodine in aqueous solutions to hydrogen bromide and iodide according to the equations:

H 3 PO 2 + 2Br 2 + 2H 2 O \u003d 4HBr + H 3 RO 4,

H 3 RO 2 + 2I 2 + 2H 2 O \u003d 4HI + H 3 RO 4.

The preparation of hypophosphites by heating white phosphorus with strong bases has been described in an experiment on the preparation of hydrogen phosphide.

When barium hypophosphite is treated with sulfuric acid, hypophosphorous acid is obtained as a result of the exchange reaction.

Everything about Red Phosphorus

PHOSPHORUS(from Greek phosphoros - luminiferous; lat. Phosphorus) - one of the most common elements of the earth's crust, located in the 3rd period, in the 5th group of the main subgroup. Its content is 0.08-0.09% of its mass. The concentration in sea water is 0.07 mg/l. It is not found in the free state due to its high chemical activity. It forms about 190 minerals, the most important of which are apatite Ca5(PO4)3(F,Cl,OH), phosphorite Ca3(PO4)2 and others. Phosphorus is found in all parts of green plants, and even more in fruits and seeds. Contained in animal tissues, is part of proteins and other essential organic compounds (ATP, DNA), is an element of life.

Story

Phosphorus discovered by the Hamburg alchemist Hennig Brand in 1669. Like other alchemists, Brand tried to find the Philosopher's Stone, but received a luminous substance. Brand focused on experiments with human urine, because he believed that it, having a golden color, may contain gold or something necessary for mining. Initially, his method consisted in the fact that at first the urine was settled for several days until it disappeared. bad smell and then boiled to a sticky state. By heating this paste to high temperatures and bringing it up to the appearance of bubbles, he hoped that, when condensed, they would contain gold. After several hours of intense boiling, grains of a white wax-like substance were obtained, which burned very brightly and, moreover, flickered in the dark. Brand named this substance phosphorus mirabilis (lat. "miraculous light carrier"). Brand's discovery of phosphorus was the first discovery of a new element since antiquity.

Somewhat later, phosphorus was obtained by another German chemist, Johann Kunkel.

Regardless of Brand and Kunkel, phosphorus was obtained by R. Boyle, who described it in the article "Method of preparing phosphorus from human urine", dated October 14, 1680 and published in 1693.

An improved method for obtaining phosphorus was published in 1743 by Andreas Marggraf.

There is evidence that Arab alchemists were able to obtain phosphorus in the 12th century.

The fact that phosphorus is a simple substance was proved by Lavoisier.

Origin of name

In 1669, Henning Brand, by heating a mixture of white sand and evaporated urine, obtained a substance glowing in the dark, first called "cold fire". The secondary name "phosphorus" comes from Greek words"φῶς" - light and "φέρω" - I carry. AT ancient Greek mythology the name Phosphorus (or Eosphorus, other Greek Φωσφόρος) was worn by the guardian of the Morning Star.

Getting Phosphorus

Phosphorus obtained from apatite or phosphorite as a result of interaction with coke and silica at a temperature of 1600 ° C:

2Ca3(PO4)2 + 10C + 6SiO2 → P4 + 10CO + 6CaSiO3

The resulting white phosphorus vapor condenses in the receiver under water. Instead of phosphorites, other compounds can be reduced, for example, metaphosphoric acid:

4HPO3 + 12C → 4P + 2H2 + 12CO

Physical Properties

Elementary phosphorus under normal conditions, it represents several stable allotropic modifications; The problem of phosphorus allotropy is complex and not fully resolved. Usually, four modifications of a simple substance are distinguished - white, red, black and metallic phosphorus. Sometimes they are also called the main allotropic modifications, implying that all the others are a variety of these four. Under normal conditions, there are only three allotropic modifications of phosphorus, and under conditions of ultrahigh pressures, there is also a metallic form. All modifications differ in color, density and other physical characteristics; there is a noticeable tendency to a sharp decrease in chemical activity during the transition from white to metallic phosphorus and an increase in metallic properties.

Red Phosphorus

Red Phosphorus, also called violet phosphorus, is a more thermodynamically stable modification of elemental phosphorus. It was first obtained in 1847 in Sweden by the Austrian chemist A. Schrötter by heating white phosphorus at 500 ° C in the atmosphere carbon monoxide(CO) in a sealed glass ampoule.

Red phosphorus has the formula Pn and is a polymer with a complex structure. Depending on the method of production and the degree of crushing of red phosphorus, it has shades from purple-red to violet, and in the cast state it has a dark purple metallic luster with a copper tint. The chemical activity of red phosphorus is much lower than that of white; it has exceptionally low solubility. It is possible to dissolve red phosphorus only in certain molten metals (lead and bismuth), which is sometimes used to obtain large crystals of it. So, for example, the German physical chemist I. V. Gittorf in 1865 for the first time received perfectly built, but small crystals (Gittorf's phosphorus). Red Phosphorus does not spontaneously ignite in air, up to a temperature of 240-250 ° C (when changing to white uniform during sublimation), but self-ignites upon friction or impact, it completely lacks the phenomenon of chemiluminescence. Insoluble in water, as well as in benzene, carbon disulfide and others, soluble in phosphorus tribromide. At the sublimation temperature, red phosphorus is converted into vapor, upon cooling of which mainly white phosphorus is formed.

Virulence Red Phosphorus thousands of times less than white, so it is used much more widely, for example, in the production of matches (the grating surface of boxes is coated with a composition based on red phosphorus)

The composition of "TERKI"

Red Phosphorus	30,8 %
Trisulfur Antimony	41,8 %
Iron Minium	12,8 %
Chalk	2,6 %
Whitewash Zinc	1,5 %
Glass ground	3,8 %
Glue Bone	6,7 %

The density of red phosphorus is also higher, reaching 2400 kg/m³ when cast. When stored in air, red phosphorus in the presence of moisture gradually oxidizes, forming a hygroscopic oxide, absorbs water and becomes damp (“soaked”), forming viscous phosphoric acid; Therefore, it is stored in an airtight container. When "soaked" - washed with water from the remnants of phosphoric acids, dried and used for its intended purpose.

Chemical properties

The chemical activity of phosphorus is much higher than that of nitrogen. The chemical properties of phosphorus are largely determined by its allotropic modification. White phosphorus is very active; in the process of transition to red and black phosphorus, the chemical activity decreases sharply. White phosphorus glows in the dark in air, the glow is due to the oxidation of phosphorus vapor to lower oxides. In the liquid and dissolved state, as well as in vapors up to 800 °C, phosphorus consists of P4 molecules. When heated above 800 °C, the molecules dissociate: Р4 = 2Р2. At temperatures above 2000 °C, molecules break up into atoms.

Interaction with Simple Substances

Phosphorus easily oxidized by oxygen:

4P + 5O2 → 2P2O5 (with excess oxygen)

4P + 3O2 → 2P2O3 (with slow oxidation or lack of oxygen)

Interacts with many simple substances - halogens, sulfur, some metals, showing oxidizing and reducing properties:

with metals - an oxidizing agent, forms phosphides:

2P + 3Ca → Ca3P2, 2P + 3Mg → Mg3P2

phosphides are decomposed by water and acids to form phosphine with non-metals - reducing agent:

2P + 3S → P2S3, 2P + 3Cl2 → 2PCl3. Does not interact with hydrogen.

Interaction with Water

Interacts with water, while disproportionate:

8P + 12H2O = 5PH3 + 3H3PO4 (phosphoric acid)

Interaction with alkalis

In alkali solutions, disproportionation occurs to a greater extent:

4P + 3KOH + 3H2O → PH3 + 3KH2PO2

Restorative Properties

Strong oxidizing agents convert phosphorus to phosphoric acid:

3P + 5HNO3 + 2H2O → 3H3PO4 + 5NO

2P + 5H2SO4 → 2H3PO4 + 5SO2 + 2H2O

The oxidation reaction also occurs when matches are ignited; Berthollet salt acts as an oxidizing agent:

6P + 5KClO3 → 5KCl + 3P2O5

Application

Phosphorus is the most important biogenic element and at the same time is very widely used in industry. Red phosphorus is used in the manufacture of matches. It, together with finely ground glass and glue, is applied to the side surface of the box. When a match head is rubbed, which includes potassium chlorate and sulfur, ignition occurs.

Toxicology of Elemental Phosphorus

red phosphorus practically non-toxic. Dust of red phosphorus, getting into the lungs, causes pneumonia with chronic action.

White phosphorus very toxic, soluble in lipids. The lethal dose of white phosphorus is 50-150 mg. Getting on the skin, white phosphorus causes severe burns.

Acute phosphorus poisoning is manifested by burning in the mouth and stomach, headache, weakness, and vomiting. After 2-3 days, jaundice develops. Chronic forms are characterized by a violation of calcium metabolism, damage to the cardiovascular and nervous systems. First aid for acute poisoning- gastric lavage, laxative, cleansing enemas, intravenous glucose solutions. In case of skin burns, treat the affected areas with solutions of copper sulfate or soda. MPC of phosphorus vapor in the air industrial premises- 0.03 mg/m³, temporary allowable concentration in atmospheric air- 0.0005 mg/m³, MPC in drinking water- 0.0001 mg/dm³.

DEFINITION

Phosphorus forms several allotropic changes: white, red and black phosphorus.

White, red and black phosphorus

White phosphorus is one of the allotropic modifications chemical element phosphorus (Fig. 1). It consists of P 4 molecules. Metastable, at room temperature soft as wax (cut with a knife), in the cold - fragile. Melts and boils without decomposition, volatile when slightly heated, distilled with water vapor. Slowly oxidizes in air (chain reaction with participation of radicals, chemiluminescence), ignites with low heating in the presence of oxygen. It dissolves well in carbon disulfide, ammonia, sulfur oxide (IV), poorly - in carbon tetrachloride. It does not dissolve in water, it is well preserved under a layer of water.

Rice. 1. White phosphorus. Appearance.

Red phosphorus is the most thermodynamically stable allotropic modification of elemental phosphorus. Under normal conditions, it is a powder of various shades (from purple-red to violet) (Fig. 2). The color is determined by the method of obtaining and the degree of crushing of the substance. Has a metallic sheen. When heated, it sublimates. Oxidizes in air. Insoluble in water and carbon disulfide. The chemical activity of red phosphorus is much less than white and black. It dissolves in a melt of lead, from which violet phosphorus (Gittorf's phosphorus) crystallizes. When red phosphorus vapor is cooled, white phosphorus is obtained.

Rice. 2. Red phosphorus. Appearance.

Black phosphorus is formed from white by heating it under high pressure at 200-220 o C. It looks like graphite, greasy to the touch. Density - 2.7 g / cm 3. Semiconductor.

Chemical formula of phosphorus

The chemical formula of white phosphorus is P 4 . It shows that the molecule of this substance contains four phosphorus atoms (Ar = 31 amu). According to the chemical formula, you can calculate molecular weight white phosphorus:

Mr(P 4) = 2×Ar(P) = 4×31 = 124.

Red phosphorus has the formula P n and is a polymer with a complex structure.

Structural (graphical) formula of phosphorus

The structural (graphical) formula of phosphorus is more visual. It shows how atoms are connected to each other within a molecule.

The structural formula of white phosphorus is:

The structural formula of the red phosphorus polymer is:

Electronic formula

An electronic formula showing the distribution of electrons in an atom over energy sublevels is shown below:

15 P 1s 2 2s 2 2p 6 3s 2 3p 3 .

It also shows that phosphorus belongs to the elements of the p-family, as well as the number of valence electrons - there are 5 electrons in the outer energy level (3s 2 3p 3).

Examples of problem solving

EXAMPLE 1

Exercise	Determine the molecular formula of a salt with a molar mass of less than 300, in which the mass fractions of nitrogen, hydrogen, chromium and oxygen are 11.11%; 3.17%; 41.27% and 44.44% respectively.
Decision	The mass fraction of the element X in the molecule of the HX composition is calculated by the following formula: ω (X) = n × Ar (X) / M (HX) × 100%. Let us denote the number of nitrogen atoms in the molecule as "x", the number of hydrogen atoms as "y", the number of chromium atoms as "z", and the number of oxygen atoms as "k". Let us find the corresponding relative atomic masses of the elements of iron and oxygen (the values ​​of the relative atomic masses taken from the Periodic Table of D.I. Mendeleev will be rounded to integers). Ar(N) = 14; Ar(H) = 1; Ar(Cr) = 52; Ar(O) = 16. We divide the percentage of elements by the corresponding relative atomic masses. Thus, we will find the relationship between the number of atoms in the molecule of the compound: x:y:z:k = m(N)/Ar(N) : m(H)/Ar(H) : m(Cr)/Ar(Cr) : m(O)/Ar(O); x:y:z:k= 11.11/14:3.17/1:41.27/52: 44.44/16; x:y:z:k= 0.79: 3.17: 0.79: 2.78 = 1: 4: 1: 3.5 = 2: 8: 2: 7. Means the simplest formula compounds of nitrogen, hydrogen, chromium and oxygen has the form N 2 H 8 Cr 2 O 7 or (NH 4) 2 Cr 2 O 7. It's ammonium dichromate.
Answer	(NH 4) 2 Cr 2 O 7

EXAMPLE 2

Exercise	As a result of combustion of oxygen-containing organic compound 1.584 g of carbon dioxide and 0.972 ml of water are collected in excess air. The vapor density of this compound in air is 1.5865. Bring out chemical formula compound if it contains two radicals of the same name.
Decision	Let's draw up a scheme for the combustion reaction of an organic compound, denoting the number of carbon, hydrogen and oxygen atoms as "x", "y" and "z", respectively: C x H y O z + O z →CO 2 + H 2 O. Let us determine the masses of the elements that make up this substance. The values ​​of relative atomic masses taken from the Periodic Table of D.I. Mendeleev, rounded up to integers: Ar(C) = 12 a.m.u., Ar(H) = 1 a.m.u., Ar(O) = 16 a.m.u. m(C) = n(C)×M(C) = n(CO 2)×M(C) = /M(C); m(H) = n(H)×M(H) = 2×n(H 2 O)×M(H) = ×M(H); m(H) =. Calculate the molar masses of carbon dioxide and water. As is known, the molar mass of a molecule is equal to the sum of the relative atomic masses of the atoms that make up the molecule (M = Mr): M(CO 2) \u003d Ar (C) + 2 × Ar (O) \u003d 12+ 2 × 16 \u003d 12 + 32 \u003d 44 g / mol; M(H 2 O) \u003d 2 × Ar (H) + Ar (O) \u003d 2 × 1 + 16 \u003d 2 + 16 \u003d 18 g / mol. m(C) = /12 = 0.432 g; m(H) = = 0.108 g. The value of the molar mass of an organic substance can be determined using its density in air: M substance = M air × D air; M substance \u003d 29 × 1.5862 \u003d 46 g / mol. Find the number of carbon and hydrogen atoms in the compound: x:y = m(C)/Ar(C) : m(H)/Ar(H); x:y = 0.432/12:0.108/1; x:y = 0.036: 0.108 = 1: 3. This means that the simplest formula of the hydrocarbon radical of this compound has the form CH 3 and molar mass 15 g/mol. This means that oxygen accounts for, which is impossible. Taking into account the condition of the problem about two radicals of the same name 2 × M (CH 3) \u003d 2 × 15 \u003d 30 g / mol, we find that oxygen accounts for, i.e. the organic oxygen-containing compound has the form CH 3 -O-CH 3 . It's acetone (dimethyl ketone).
Answer	CH 3 -O-CH 3

Phosphorus is known in several allotropic modifications: white, red, violet and black. In laboratory practice, one has to meet with white and red modifications.

White phosphorus is a solid. Under normal conditions, it is yellowish, soft and similar in appearance to wax. It is easily oxidized and flammable. White phosphorus is poisonous - it leaves painful burns on the skin. White phosphorus goes on sale in the form of sticks of different lengths with a diameter of 0.5-2 cm.

White phosphorus is easily oxidized, and therefore it is stored under water in carefully sealed dark glass vessels in poorly lit and not very cold rooms (to avoid cracking the jars due to freezing water). The amount of oxygen contained in water and oxidizing phosphorus is very small; it is 7-14 mg per liter of water.

Under the influence of light, white phosphorus turns into red.

With slow oxidation, the glow of white phosphorus is observed, and with vigorous oxidation, it ignites.

White phosphorus is taken with tweezers or metal tongs; in no case should you touch it with your hands.

In case of a burn with white phosphorus, the burned area is washed with a solution of AgNO 3 (1:1) or KMnO 4 (1:10) and a wet dressing is applied soaked in the same solutions or a 5% solution of copper sulfate, then the wound is washed with water and after smoothing the epidermis, apply vaseline bandage with methyl violet. For severe burns, see a doctor.

Solutions of silver nitrate, potassium permanganate and copper sulfate oxidize white phosphorus and thus stop its damaging effect.

In case of white phosphorus poisoning, take a teaspoon of a 2% copper sulfate solution orally until vomiting occurs. Then, using the Mitcherlich test, based on the luminescence, the presence of phosphorus is determined. For this, water acidified with sulfuric acid is added to the vomiting of the poisoned person, and distilled in the dark; at the content of phosphorus, the glow of vapors is observed. A Wurtz flask is used as a device, to the side tube of which a Liebig condenser is attached, from where the distilled products enter the receiver. If phosphorus vapor is directed into a solution of silver nitrate, then a black precipitate of metallic silver forms, which is formed according to the equation given in the experiment on the reduction of silver salts with white phosphorus.

Already 0.1 G white phosphorus is a lethal dose for an adult.

White phosphorus is cut with a knife or scissors in a porcelain mortar under water. When using water at room temperature, phosphorus crumbles. Therefore, it is better to use warm water, but not higher than 25-30 °. After cutting the phosphorus in warm water, it is transferred to cold water or cooled with a stream of cold water.

White phosphorus is a highly flammable substance. It ignites at a temperature of 36-60°, depending on the concentration of oxygen in the air. Therefore, when conducting experiments, in order to avoid an accident, it is necessary to take into account every grain of it.

Drying of white phosphorus is carried out by quickly applying thin asbestos or filter paper to it, avoiding friction or pressure.

When phosphorus ignites, it is extinguished with sand, a wet towel or water. If burning phosphorus is on a sheet of paper (or asbestos), this sheet must not be touched, since molten burning phosphorus can be easily spilled.

White phosphorus melts at 44°, boils at 281°. White phosphorus is melted with water, because in contact with air, molten phosphorus ignites. By fusion and subsequent cooling, white phosphorus can be easily recovered from the waste. To do this, white phosphorus waste from various experiments, collected in a porcelain crucible with water, is heated in a water bath. If crust formation is noticeable on the surface of the molten phosphorus, a little HNO 3 or a chromium mixture is added. The crust is oxidized, small grains merge into a common mass, and after cooling with a jet of cold water, one piece of white phosphorus is obtained.

Under no circumstances should phosphorus residues be thrown into the sink, since, accumulating in the bends of the elbow of sewers, it can cause burns to maintenance workers.

Experience. Melting and supercooling of molten white phosphorus. A piece of white phosphorus the size of a pea is placed in a test tube with water. The test tube is placed in a beaker filled almost to the top with water and fixed in a vertical position in a tripod clamp. The glass is slightly heated and using a thermometer determine the temperature of the water in the test tube at which phosphorus melts. After the end of melting, the tube is transferred to a beaker with cold water and solidification of phosphorus is observed. If the tube is stationary, then at a temperature below 44° (up to 30°) white phosphorus remains in a liquid state.

The liquid state of white phosphorus, cooled below its melting point, is a state of supercooling.

After the end of the experiment, in order to more easily extract phosphorus, it is melted again and the test tube is immersed with the hole up in an inclined position in a vessel with cold water.

Experience. Attaching a piece of white phosphorus to the end of the wire. To melt and solidify white phosphorus, a small porcelain crucible with phosphorus and water is used; it is placed in a glass of warm and then cold water. The wire for this purpose is taken iron or copper with a length of 25-30 cm and diameter 0.1-0.3 cm. When the wire is immersed in solidifying phosphorus, it easily attaches to it. In the absence of a crucible, a test tube is used. However, due to the insufficiently even surface of the test tube, it is sometimes necessary to break it in order to extract the phosphorus. To remove white phosphorus from the wire, it is immersed in a glass of warm water.

Experience. Determination of the specific gravity of phosphorus. At 10°, the specific gravity of phosphorus is 1.83. Experience allows us to make sure that white phosphorus is heavier than water and lighter than concentrated H 2 SO 4.

When a small piece of white phosphorus is introduced into a test tube with water and concentrated H 2 SO 4 (specific weight 1.84), it is observed that phosphorus sinks in water, but floats on the surface of the acid, melting due to the heat released when concentrated H 2 SO is dissolved 4 in water.

To pour concentrated H 2 SO 4 into a test tube with water, use a funnel with a long and narrow neck, reaching to the end of the test tube. Pour in the acid and remove the funnel from the test tube carefully so as not to cause mixing of the liquids.

At the end of the experiment, the contents of the test tube are stirred with a glass rod and cooled from the outside with a stream of cold water until the phosphorus solidifies so that it can be removed from the test tube.

When using red phosphorus, it is observed that it sinks not only in water, but also in concentrated H 2 SO 4, since its specific gravity (2.35) is greater than the specific gravity of both water and concentrated sulfuric acid.

WHITE PHOSPHORUS, GLOW

Due to the slow oxidation that occurs even at ordinary temperatures, white phosphorus glows in the dark (hence the name "luminiferous"). Around a piece of phosphorus in the dark, a greenish luminous cloud appears, which, when the phosphorus vibrates, is set in a wave-like motion.

Phosphorescence (luminescence of phosphorus) is explained by the slow oxidation of phosphorus vapor by oxygen in the air to phosphorous and phosphorus anhydride with the release of light, but without the release of heat. In this case, ozone is released, and the air around is ionized (see the experiment showing the slow combustion of white phosphorus).

Phosphorescence depends on temperature and oxygen concentration. At 10° and normal pressure phosphorescence proceeds weakly, and in the absence of air does not occur at all.

Substances that react with ozone (H 2 S, SO 2, Cl 2, NH 3, C 2 H 4, turpentine oil) weaken or completely stop phosphorescence.

The conversion of chemical energy into light energy is called "chemiluminescence".

Experience. Observation of the glow of white phosphorus. If you observe in the dark a piece of white phosphorus in a glass and not completely covered with water, then a greenish glow is noticeable. In this case, wet phosphorus slowly oxidizes, but does not ignite, since the temperature of the water is below the flash point of white phosphorus.

The glow of white phosphorus can be observed after a piece of white phosphorus has been exposed to air for a short time. If you put a few pieces of white phosphorus in a flask on glass wool and fill the flask with carbon dioxide, lowering the end of the outlet tube to the bottom of the flask under glass wool, and then slightly heat the flask by dipping it into a vessel with warm water, then in the dark you can observe the formation of a cold pale greenish flame (you can safely put your hand into it).

The formation of a cold flame is explained by the fact that carbon dioxide leaving the flask entrains phosphorus vapor, which begin to oxidize when it comes into contact with air at the opening of the flask. In a flask, white phosphorus does not ignite, because it is in an atmosphere of carbon dioxide. At the end of the experiment, the flask is filled with water.

In describing the experiment for obtaining white phosphorus in an atmosphere of hydrogen or carbon dioxide, it was already mentioned that carrying out these experiments in the dark makes it possible to observe the glow of white phosphorus.

If you make an inscription on a wall, a sheet of cardboard or paper with phosphor chalk, then due to phosphorescence the inscription remains visible for a long time in the dark.

Such an inscription cannot be made on a blackboard, since after that ordinary chalk does not stick to it and the board has to be washed with gasoline or another stearin solvent.

Phosphorus chalk is obtained by dissolving liquid white phosphorus in molten stearin or paraffin. To do this, approximately two parts by weight of stearin (candle pieces) or paraffin are added to a test tube to one weight part of dry white phosphorus, the test tube is covered with cotton wool to prevent oxygen from entering, and heated with continuous shaking. After the end of melting, the test tube is cooled with a jet of cold water, then the test tube is broken and the solidified mass is removed.

Phosphorus chalk is stored under water. When using a piece of such chalk is wrapped in wet paper.

Phosphorus chalk can also be obtained by adding small pieces of dried white phosphorus to paraffin (stearin) melted in a porcelain cup. If the paraffin ignites when phosphorus is added, it is extinguished by covering the cup with a piece of cardboard or asbestos.

After some cooling, the solution of phosphorus in paraffin is poured into dry and clean test tubes and cooled with a stream of cold water until it solidifies into a solid mass.

After that, test tubes are broken, chalk is removed and stored under water.

SOLUBILITY OF WHITE PHOSPHORUS

In water, white phosphorus is sparingly soluble, slightly soluble in alcohol, ether, benzene, xylene, methyl iodide and glycerin; dissolves well in carbon disulfide, sulfur chloride, phosphorus trichloride and tribromide, carbon tetrachloride.

Experience. Dissolution of white phosphorus in carbon disulfide. Carbon disulfide is a colorless, highly volatile, highly flammable, poisonous liquid. Therefore, when working with it, avoid inhaling its vapors and turn off all gas burners.

Three or four pieces of white phosphorus the size of a pea are dissolved with light shaking in a glass of 10-15 ml carbon disulfide.

If a small sheet of filter paper is moistened with this solution and held in air, the paper ignites after a while. This is because carbon disulfide evaporates quickly, and the finely divided white phosphorus remaining on the paper quickly oxidizes at ordinary temperatures and ignites due to the heat released during oxidation. (It is known that the ignition temperature of various substances depends on the degree of their grinding.) It happens that paper does not ignite, but only chars. Paper moistened with a solution of phosphorus in carbon disulfide is kept in air with metal tongs.

The experiment is carried out carefully so that drops of a solution of phosphorus in carbon disulfide do not fall on the floor, on the table, on clothes or on hands.

If the solution gets on the hand, it is quickly washed with soap and water, and then with a solution of KMnO 4 (to oxidize particles of white phosphorus that have fallen on the hands).

The solution of phosphorus in carbon disulfide remaining after the experiments is not stored in the laboratory, since it can easily ignite.

WHITE PHOSPHORUS TRANSFORMATION TO RED

White phosphorus is converted to red according to the equation:

P (white) = P (red) + 4 kcal.

Installation for the production of white phosphorus from red: test tube-reactor 1, tube 2, through which carbon dioxide enters the test tube-reactor, gas outlet tube 3, through which vapors of white phosphorus, together with carbon dioxide, leave the test tube and are cooled with water

The process of converting white phosphorus to red is greatly accelerated by heating, under the influence of light and in the presence of traces of iodine (1 G iodine at 400 G white phosphorus). Iodine, combining with phosphorus, forms phosphorus iodide, in which white phosphorus dissolves and quickly turns into red with the release of heat.

Red phosphorus is obtained by prolonged heating of white phosphorus in a closed vessel in the presence of traces of iodine to 280-340 °

With long-term storage of white phosphorus in the light, it gradually turns into red.

Experience. Obtaining a small amount of red phosphorus from white. In a glass tube 10-12 long, closed at one end cm and diameter 0.6-0.8 cm they introduce a piece of white phosphorus the size of a grain of wheat and a very small crystal of iodine. The tube is sealed and suspended in an air bath over a tray of sand, then heated to 280-340° and the transformation of white phosphorus into red is observed.

Partial conversion of white phosphorus to red can also be observed by slightly heating a test tube with a small piece of white phosphorus and a very small crystal of iodine. Before starting heating, the test tube is closed with a swab of glass (asbestos or ordinary) wool and a tray with sand is placed under the test tube. The tube is heated for 10-15 minutes (without bringing the phosphorus to a boil) and the transformation of white phosphorus into red is observed.

White phosphorus remaining in the test tube can be removed by heating with a concentrated alkali solution or by burning.

The transformation of white phosphorus into red can also be observed by heating a small piece of phosphorus in a test tube in an atmosphere of carbon dioxide to a temperature below boiling.

COMBUSTION OF WHITE PHOSPHORUS

When white phosphorus burns, phosphoric anhydride is formed:

P 4 + 5O 2 \u003d 2P 2 O 5 + 2 x 358.4 kcal.

You can observe the combustion of phosphorus in air (slow and fast) and under water.

Experience. Slow combustion of white phosphorus and air composition. This experiment was not described as a way to obtain nitrogen, since it does not completely bind the oxygen contained in the air.

The slow oxidation of white phosphorus by atmospheric oxygen occurs in two stages; in the first stage, phosphorous anhydride and ozone are formed according to the equations:

2P + 2O 2 \u003d P 2 O 3 + O, O + O 2 \u003d O 3.

In the second stage, phosphorous anhydride is oxidized to phosphoric anhydride.

The slow oxidation of white phosphorus is accompanied by luminescence and ionization of the surrounding air.

An experiment showing the slow burning of white phosphorus should last at least three hours. The apparatus required for the experiment is shown in Fig.

In a cylinder expanded at the opening, almost filled with water, a graduated tube with a closed end, containing about 10 ml water. Tube length 70 cm, diameter 1.5-2 cm. After lowering the graduated tube, remove the finger from the tube opening, bring the water in the tube and cylinder to the same level, and note the volume of air contained in the tube. Without raising the tube above the water level in the cylinder (in order not to let in additional air), a piece of white phosphorus fixed at the end of the wire is introduced into the air space of the tube.

After three to four hours, or even after two or three days, a rise in water in the tube is noted.

At the end of the experiment, the wire with phosphorus is removed from the tube (without raising the tube above the water level in the cylinder), the water in the tube and cylinder is brought to the same level, and the volume of air remaining after the slow oxidation of white phosphorus is noted.

Experience shows that as a result of the binding of oxygen by phosphorus, the volume of air has decreased by one fifth, which corresponds to the oxygen content in the air.

Experience. Rapid combustion of white phosphorus. Due to the fact that a large amount of heat is released during the reaction of the combination of phosphorus with oxygen, white phosphorus ignites spontaneously in air and burns with a bright yellowish-white flame, forming phosphorus anhydride, a white solid that combines very vigorously with water.

It was already mentioned earlier that white phosphorus ignites at 36-60 °. To observe its self-ignition and combustion, a piece of white phosphorus is placed on a sheet of asbestos and covered with a glass bell or a large funnel, on the neck of which a test tube is put on.

Phosphorus can be easily set on fire with a glass rod heated in hot water.

Experience. Comparison of ignition temperatures of white and red phosphorus. At one end of a copper plate (length 25 cm, width 2.5 cm and thickness 1 mm) put a small piece of dried white phosphorus, pour a small pile of red phosphorus on the other end. The plate is placed on a tripod and at the same time approximately equally burning gas burners are brought to both ends of the plate.

White phosphorus ignites immediately, and red phosphorus only when its temperature reaches approximately 240°.

Experience. Ignition of white phosphorus under water. A test tube with water containing several small pieces of white phosphorus is dipped into a glass of hot water. When the water in the test tube is heated to 30-50°C, a current of oxygen is passed into it through the tube. Phosphorus ignites and burns, scattering bright sparks.

If the experiment is carried out in the beaker itself (without a test tube), the beaker is placed on a tripod mounted on a tray of sand.

REDUCTION OF SILVER AND COPPER SALTS WITH WHITE PHOSPHORUS

Experience. When a piece of white phosphorus is introduced into a test tube with a solution of silver nitrate, a precipitate of metallic silver is observed (white phosphorus is an energetic reducing agent):

P + 5AgNO 3 + 4H 2 O \u003d H 3 RO 4 + 5Ag + 5HNO 3.

If white phosphorus is introduced into a test tube with a solution of copper sulfate, then metallic copper precipitates:

2P + 5CuSO 4 + 8H 2 O \u003d 2H 3 PO 4 + 5H 2 SO 4 + 5Cu.