Production and production of phosphorus. Getting white phosphorus

Phosphorus and its compounds

Introduction

Chapter I. Phosphorus as an element and as a simple substance

1.1. Phosphorus in nature

1.2. Physical Properties

1.3. Chemical properties

1.4. Receipt

1.5. Application

Chapter II. Phosphorus compounds

2.1. oxides

2.2. Acids and their salts

2.3. Phosphine

Chapter III. Phosphate fertilizers

Conclusion

Bibliographic list

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. outdoor energy level A phosphorus atom has five electrons. Graphically it looks like this:

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

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

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.

Chapter I Phosphorus as an element and as a simple substance

1.1 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 - necessary element living organisms. It is present in bones, muscles, brain tissue and nerves. ATP molecules are built from phosphorus - adenosine tri phosphoric acid(ATP is a 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.

1.2 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 linked by three σ-bonds to the other three atoms.

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

red phosphorus- a 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 By appearance similar to graphite, greasy to the touch, has semiconductor properties. Density 2.7 g/cm 3 .

Red and black phosphorus have an atomic crystal lattice.

1.3 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 \u003d 2P 2 O 5 (phosphorus oxide (V), phosphoric anhydride).

II. Interaction with complex substances.

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

1.4 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

6CaSiO3 + P4 + 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).

1.5 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.

KCl + .

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 metal 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.

Chapter II . Phosphorus compounds

2.1 Oxides

Phosphorus forms several oxides. The most important of these are phosphorus oxide (V) P 4 O 10 and phosphorus oxide (III) P 4 O 6 . Often their formulas are written in a simplified form - P 2 O 5 and P 2 O 3. The structure of these oxides retains the tetrahedral arrangement of phosphorus atoms.

Phosphorus oxide(III) P 4 O 6 is a waxy crystalline mass that melts at 22.5°C and turns into a colorless liquid. Poisonous.

When dissolved in cold water forms phosphorous acid:

P 4 O 6 + 6H 2 O \u003d 4H 3 PO 3,

and when reacting with alkalis, the corresponding salts (phosphites).

Strong reducing agent. When interacting with oxygen, it is oxidized to P 4 O 10.

Phosphorus (III) oxide is obtained by the oxidation of white phosphorus in the absence of oxygen.

Phosphorus oxide(V) P 4 O 10 - white crystalline powder. The sublimation temperature is 36°C. It has several modifications, one of which (the so-called volatile) has the composition P 4 O 10 . The crystal lattice of this modification is composed of P 4 O 10 molecules interconnected by weak intermolecular forces, which are easily broken when heated. Hence the volatility of this variety. Other modifications are polymeric. They are formed by infinite layers of PO 4 tetrahedra.

When P 4 O 10 interacts with water, phosphoric acid is formed:

P 4 O 10 + 6H 2 O \u003d 4H 3 PO 4.

Being acid oxide, P 4 O 10 reacts with basic oxides and hydroxides.

It is formed during high-temperature oxidation of phosphorus in excess oxygen (dry air).

Due to its exceptional hygroscopicity, phosphorus (V) oxide is used in laboratory and industrial technology as a drying and dehydrating agent. In its drying effect, it surpasses all other substances. Chemically bound water is removed from anhydrous perchloric acid with the formation of its anhydride:

4HClO 4 + P 4 O 10 \u003d (HPO 3) 4 + 2Cl 2 O 7.

2.2 Acids and their salts

A) Phosphorous acid H3PO3. Anhydrous phosphorous acid H 3 PO 3 forms crystals with a density of 1.65 g/cm 3 , melting at 74°C.

Structural formula:

When anhydrous H 3 RO 3 is heated, a disproportionation reaction (self-oxidation-self-recovery) occurs:

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

Salts of phosphorous acid - phosphites. For example, K 3 PO 3 (potassium phosphite) or Mg 3 (PO 3) 2 (magnesium phosphite).

Phosphorous acid H 3 RO 3 is obtained by dissolving phosphorus (III) oxide in water or by hydrolysis of phosphorus (III) chloride РCl 3:

РCl 3 + 3H 2 O \u003d H 3 PO 3 + 3HCl.

b) Phosphoric acid (orthophosphoric acid)H3PO4.

Anhydrous phosphoric acid is a light transparent crystals, with room temperature floating in the air. Melting point 42.35°C. With water, phosphoric acid forms solutions of any concentration.

Designation - P (Phosphorus);
Period - III;
Group - 15 (Va);
Atomic mass - 30.973761;
Atomic number - 15;
Radius of an atom = 128 pm;
Covalent radius = 106 pm;
Electron distribution - 1s 2 2s 2 2p 6 3s 2 3p 3 ;
melt t = 44.14°C;
boiling point = 280°C;
Electronegativity (according to Pauling / according to Alpred and Rochov) = 2.19 / 2.06;
Oxidation state: +5, +3, +1, 0, -1, -3;
Density (n.a.) \u003d 1.82 g / cm 3 (white phosphorus);
Molar volume = 17.0 cm 3 / mol.

Phosphorus compounds:

Phosphorus (carrying light) was first obtained by the Arab alchemist Ahad Behil in the 12th century. Of the European scientists, the German Hennig Brant was the first to discover phosphorus in 1669, during experiments with human urine in an attempt to extract gold from it (the scientist believed that the golden color of urine was caused by the presence of gold particles). Somewhat later, phosphorus was obtained by I. Kunkel and R. Boyle - the latter described it in his article "Method of preparing phosphorus from human urine" (10/14/1680; the work was published in 1693). Lavoisier later proved that phosphorus is a simple substance.

The content of phosphorus in the earth's crust is 0.08% by mass - this is one of the most common chemical elements on our planet. Due to its high activity, phosphorus in a free state does not occur in nature, but is part of almost 200 minerals, the most common of which are Ca 5 (PO 4) 3 (OH) apatite and Ca 3 (PO 4) 2 phosphorite.

Phosphorus plays an important role in the life of animals, plants and humans - it is part of such a biological compound as a phospholipid, it is also present in protein and other such important organic compounds like DNA and ATP.

Rice. The structure of the phosphorus atom.

The phosphorus atom contains 15 electrons, and has a similar to nitrogen electronic configuration external valence level (3s 2 3p 3), but phosphorus has less pronounced non-metallic properties than nitrogen, which is explained by the presence of a free d-orbital, a large atomic radius and a lower ionization energy.

Entering into reactions with other chemical elements, the phosphorus atom can show an oxidation state from +5 to -3 (the most typical oxidation state is +5, the rest are quite rare).

+5 - phosphorus oxide P 2 O 5 (V); phosphoric acid (H 3 PO 4); phosphates, halides, sulfides of phosphorus V (salts of phosphoric acid);
+3 - P 2 O 3 (III); phosphorous acid (H 3 PO 3); phosphites, halides, sulfides of phosphorus III (salts of phosphorous acid);
0-P;
-3 - phosphine PH 3; metal phosphides.

In the ground (unexcited) state, the phosphorus atom has two paired electrons in the s-sublevel + 3 unpaired electrons in the p-orbitals (the d-orbital is free) at the outer energy level. In the excited state, one electron from the s-sublevel passes to the d-orbital, which expands the valence possibilities of the phosphorus atom.

Rice. The transition of the phosphorus atom to an excited state.

P2

Two phosphorus atoms are combined into a P 2 molecule at a temperature of about 1000°C.

With more low temperatures phosphorus exists in the four-atom P 4 molecules, as well as in the more stable polymer molecules P ∞ .

Allotropic modifications of phosphorus:

White phosphorus- extremely poisonous (the lethal dose of white phosphorus for an adult is 0.05-0.15 g) waxy substance with the smell of garlic, without color, luminous in the dark (slow oxidation process in P 4 O 6); the high reactivity of white phosphorus is explained by weak P-P connections(white phosphorus has a molecular crystal lattice with the formula P 4, at the nodes of which phosphorus atoms are located), which are quite easily broken, as a result of which white phosphorus, when heated or during long-term storage, passes into more stable polymer modifications: red and black phosphorus. For these reasons, white phosphorus is stored without air access under a layer of purified water or in special inert media.
yellow phosphorus- a flammable, highly toxic substance, does not dissolve in water, easily oxidizes in air and ignites spontaneously, while burning with a bright green dazzling flame with the release of thick white smoke.
red phosphorus- a polymeric, water-insoluble substance with a complex structure, which has the least reactivity. Red phosphorus is widely used in industrial production, because it is not so toxic. Since in the open air, red phosphorus, absorbing moisture, gradually oxidizes to form a hygroscopic oxide (“damp”), forms viscous phosphoric acid, therefore, red phosphorus is stored in a hermetically sealed container. In the case of soaking, red phosphorus is purified from phosphoric acid residues by washing with water, then dried and used for its intended purpose.
black phosphorus- greasy to the touch graphite-like substance of gray-black color, with semiconductor properties - the most stable modification of phosphorus with an average reactivity.
metallic phosphorus obtained from black phosphorus high pressure. Metallic phosphorus conducts electricity very well.

Chemical properties of phosphorus

Of all the allotropic modifications of phosphorus, the most active is white phosphorus (P 4). Often in the equation chemical reactions just write P, not P 4 . Since phosphorus, like nitrogen, has many variants of oxidation states, in some reactions it is an oxidizing agent, in others it is a reducing agent, depending on the substances with which it interacts.

Oxidative phosphorus exhibits properties in reactions with metals that occur when heated to form phosphides:
3Mg + 2P \u003d Mg 3 P 2.

Phosphorus is reducing agent in reactions:

with more electronegative non-metals (oxygen, sulfur, halogens):
- phosphorus (III) compounds are formed with a lack of an oxidizing agent
  4P + 3O 2 \u003d 2P 2 O 3
- phosphorus (V) compounds - with an excess of: oxygen (air)
  4P + 5O 2 \u003d 2P 2 O 5
with halogens and sulfur, phosphorus forms halides and sulfide of 3- or 5-valent phosphorus, depending on the ratio of reagents, which are taken in deficiency or excess:
- 2P + 3Cl 2 (week) \u003d 2PCl 3 - phosphorus (III) chloride
- 2P + 3S (weeks) \u003d P 2 S 3 - phosphorus (III) sulfide
- 2P + 5Cl2 (ex.) \u003d 2PCl 5 - phosphorus (V) chloride
- 2P + 5S (ex.) \u003d P 2 S 5 - phosphorus (V) sulfide
with concentrated sulfuric acid:
2P + 5H 2 SO 4 \u003d 2H 3 PO 4 + 5SO 2 + 2H 2 O
with concentrated nitric acid:
P + 5HNO 3 \u003d H 3 PO 4 + 5NO 2 + H 2 O
with diluted nitric acid:
3P + 5HNO 3 + 2H 2 O \u003d 3H 3 PO 4 + 5NO

Phosphorus acts as both an oxidizing agent and a reducing agent in reactions disproportionation With aqueous solutions alkalis when heated, forming (except for phosphine) hypophosphites (salts of hypophosphorous acid), in which it exhibits an uncharacteristic oxidation state +1:
4P 0 + 3KOH + 3H 2 O \u003d P -3 H 3 + 3KH 2 P +1 O 2

REMEMBER: with other acids, except for the above reactions, phosphorus does not react.

Getting and using phosphorus

Industrially, phosphorus is produced by its reduction with coke from phosphorites (fluorapatates), which include calcium phosphate, by calcining in electric furnaces at a temperature of 1600 ° C with the addition of quartz sand:
Ca 3 (PO 4) 2 + 5C + 3SiO 2 = 3CaSiO 3 + 2P + 5CO.

At the first stage of the reaction, under the influence of high temperature, silicon (IV) oxide displaces phosphorus (V) oxide from phosphate:
Ca 3 (PO 4) 2 + 3SiO 2 \u003d 3CaSiO 3 + P 2 O 5.

Then phosphorus oxide (V) is reduced by coal to free phosphorus:
P 2 O 5 + 5C \u003d 2P + 5CO.

The use of phosphorus:

pesticides;
matches;
detergents;
paints;
semiconductors.

PHOSPHORUS, P (lat. Phosphorus * a. phosphorus; n. Phosphor; f. phosphore; and. fosforo), is a chemical element of group V of the periodic system of Mendeleev, atomic number 15, atomic mass 30.97376. Natural phosphorus is represented by one stable isotope 31 P. Six artificial radioactive isotopes of phosphorus are known with mass numbers 28-30 and 32-34.

The method of obtaining phosphorus may have been known to Arab alchemists as early as the 12th century, but the generally accepted date for the discovery of phosphorus is 1669, when H. Brand () received a substance that glows in the dark, called "cold fire". The existence of phosphorus chemical element proved in the early 1970s. 18th century French chemist A. Lavoisier.

Modifications and properties

Elemental phosphorus exists in the form of several allotropic modifications - white, red, black. White phosphorus is a waxy transparent substance with a characteristic odor, formed by the condensation of phosphorus vapor. In the presence of impurities - traces of red phosphorus, arsenic, iron, etc. - colored in yellow, so commercial white phosphorus is called yellow. There are 2 versions of white phosphorus a-P has a cubic lattice of the closest packing a=0.185 nm; density 1828 kg/m3; melting point 44.2°C, boiling point 277°C; thermal conductivity 0.56 W/(m.K); molar heat capacity 23.82 J / (mol.K); temperature coefficient of linear expansion 125.10 -6 K -1 ; in electrical properties, white phosphorus is close to dielectrics. At a temperature of 77.8 ° C and a pressure of 0.1 MPa, a-P turns into b-P (rhombic lattice, density 1880 kg / m 3). Heating white phosphorus in the absence of air at 250-300°C for several hours leads to the formation of a red modification. Ordinary commercial red phosphorus is practically amorphous, however, with prolonged heating, it can turn into one of the crystalline forms (triclinic, cubic) with a density of 2000 to 2400 kg / m 3 and a melting point of 585-610 ° C. During sublimation (sublimation t 431 ° C), red phosphorus turns into a gas, upon cooling of which mainly white phosphorus is formed. When white phosphorus is heated to 200–220°C under a pressure of 1.2–1.7 GPa, black phosphorus is formed. This type transformations can be carried out with normal pressure(at t 370°C), using as a catalyst, as well as a small amount of black phosphorus for seeding. Black phosphorus is a crystalline substance with a rhombic lattice (a=0.331, b=0.438 and c=1.05 nm), density 2690 kg/m 3 , melting point 1000 °C; similar in appearance to graphite; semiconductor, diamagnetic. When heated to a temperature of 560-580°C and pressure saturated vapors turns into red phosphorus.

Chemical phosphorus

Phosphorus atoms are combined into diatomic (P 2) and tetraatomic (P 4) polymer molecules. The most stable under normal conditions are molecules containing long chains of interconnected P4 tetrahedra. In compounds, phosphorus has an oxidation state of +5, +3, -3. Like nitrogen in chemical compounds, it forms mainly covalent bond. Phosphorus is a reactive element. Its white modification is most active, which ignites spontaneously at a temperature of about 40 ° C, therefore it is stored under a layer of water. Red phosphorus ignites on impact or friction. Black phosphorus is inactive and hardly ignites when ignited. The oxidation of phosphorus is usually accompanied by chemiluminescence. During the combustion of phosphorus in an excess of oxygen, P 2 O 5 is formed, with a deficiency, mainly P 2 O 3. Phosphorus forms acids: ortho- (H 3 PO 4), polyphosphoric (H n + 2 PO 3n + 1), phosphorous (H 3 PO 3), phosphorous (H 4 P 2 O 6), phosphorous (H 3 PO 2) , as well as peracids: perphosphoric (H 4 P 2 O 8) and monoperphosphoric (H 3 PO 5).

Phosphorus reacts directly with all halogens, releasing large amounts of heat. Phosphorus sulfides and nitrides are known. At a temperature of 2000°C, phosphorus reacts with carbon to form carbide (PC 3); when phosphorus is heated with metals - phosphides. White phosphorus and its compounds are highly toxic, MPC 0.03 mg/m 3 .

Phosphorus in nature

The average content of phosphorus in the earth's crust (clarke) is 9.3.10 -2%, in ultrabasic rocks 1.7. 10 -2%, basic - 1.4.10 -2%, acidic - 7.10 -2%, sedimentary - 7.7.10 -2%. Phosphorus is involved in magmatic processes and migrates vigorously in the biosphere. Both processes are associated with its large accumulations, which form industrial deposits of apatites - Ca 5 (PO 4) 3 (F, Cl) and phosphorites - amorphous Ca 5 (PO 4) 3 (OH, CO 3) with various impurities. Phosphorus is an extremely important biogenic element that is accumulated by many organisms. It is with biogenic migration that the processes of phosphorus concentration in the earth's crust are associated. Over 180 minerals containing phosphorus are known.

Getting and using

On an industrial scale, phosphorus is extracted from natural phosphates by electrothermal reduction with coke at temperatures of 1400–1600°C in the presence of silica (quartz sand); gaseous phosphorus after cleaning from dust is sent to condensing units, where liquid technical white phosphorus is collected under a layer of water. The bulk of the produced phosphorus is processed into phosphoric acid and obtained based on it. phosphate fertilizers and technical salts. Salts of phosphoric acids are widely used - phosphates, to a somewhat lesser extent - phosphites and hypophosphites. White phosphorus is used in the manufacture of incendiary and smoke projectiles; red - in match production.

Phosphorus is one of the fairly common elements; its content in the earth's crust is about . Due to the easy oxidizability of phosphorus in the free state, it does not occur in nature.

Of the natural compounds of phosphorus, the most important is calcium orthophosphate, which sometimes forms large deposits in the form of the mineral phosphorite. In the USSR, the richest deposits of phosphorites are located in southern Kazakhstan in the Karatau mountains. The mineral apatite is also often found, containing, in addition to, more or. Huge deposits of apatite were discovered in the twenties of our century on the Kola Peninsula.

This deposit is the largest in the world in terms of its reserves.

Phosphorus, like nitrogen, is necessary for all living beings, as it is part of some proteins of both plant and animal origin. In plants, phosphorus is found mainly in seed proteins, in animal organisms - in milk, blood, cerebral and nervous proteins. In addition, a large amount of phosphorus is found in the bones of vertebrates, mainly in the form of compounds and. In the form of an acid residue of phosphoric acid, phosphorus is part of nucleic acids - complex organic polymer compounds found in all living organisms. These acids are directly involved in the transfer of hereditary properties of a living cell.

The raw materials for the production of phosphorus and its compounds are phosphorites and apatites. Natural phosphorite or apatite is crushed, mixed with sand and coal and heated in furnaces using electric current without air access.

To understand the reaction taking place, imagine calcium phosphate as a compound of calcium oxide with phosphoric anhydride, while sand consists mainly of silicon dioxide. At high temperature silicon dioxide displaces phosphorus anhydride and, combining with calcium oxide, forms low-melting calcium silicate, and phosphorus anhydride is reduced by coal to free phosphorus:

Adding both equations, we get:

Phosphorus is released in the form of vapors, which condense in the receiver underwater.

Phosphorus forms several allotropic modifications.

White phosphorus is obtained in the solid state when rapid cooling phosphorus vapor; its density. In its pure form, white phosphorus is completely colorless and transparent; the sales product is usually colored yellowish and waxy in appearance. In the cold, white phosphorus is brittle, but at higher temperatures it becomes soft and can be easily cut with a knife.

In air, white phosphorus oxidizes very quickly and glows in the dark. Hence the name "phosphorus", which in Greek means "light-bearing". Already at low heating, for which simple friction is sufficient, phosphorus burns out, releasing a large amount of heat. Phosphorus can also spontaneously ignite in air due to the release of heat during oxidation.

To protect white phosphorus from oxidation, it is stored under water. White phosphorus is insoluble in water; readily soluble in carbon disulfide.

White phosphorus has a molecular crystal lattice, at the nodes of which are tetrahedral molecules. The bond strength between atoms in these molecules is relatively small. This explains the high chemical activity of white phosphorus.

White phosphorus is a strong poison, even in small doses it is deadly.

If white phosphorus is heated for a long time without access to air at , then it turns into another modification of phosphorus, which has a red-violet color and is called red phosphorus. The same transformation occurs, but only very slowly, under the action of light.

Red phosphorus differs sharply from white in its properties: it oxidizes very slowly in air, does not glow in the dark, lights up only at , does not dissolve in carbon disulfide and is non-toxic. The density of red phosphorus is . The variable density value is due to the fact that red phosphorus consists of several forms. Their structure is not fully elucidated, but it is known that they are polymeric substances.

With strong heating, red phosphorus, without melting, evaporates (sublimes). When the vapor is cooled, white phosphorus is obtained.

Black phosphorus is formed from white when heated to a very high pressure. It looks like graphite, is greasy to the touch and heavier than other modifications; its density is . Black phosphorus is a semiconductor.

The use of phosphorus is very diverse. A large amount of it is spent on the production of matches.

In the manufacture of matches, red phosphorus is used; it is contained in the mass, which is applied to matchbox. The head of a match consists of a mixture of combustible substances with salt and compounds that catalyze the decomposition of salt, etc.)

In addition to match production, phosphorus is used in metallurgy. It is used to make some semiconductors, gallium phosphide, indium phosphide. It is introduced into the composition of other semiconductors in very small quantities as a necessary additive. In addition, it is part of some metallic materials, such as tin bronzes.

When burning phosphorus, thick white smoke is produced; therefore, white phosphorus is used to equip ammunition (artillery shells, air bombs, etc.) intended to form smoke screens.

A large amount of phosphorus is used in the production of organophosphorus preparations, which include very effective means destruction of insect pests.

Free phosphorus is extremely active. It directly interacts with many simple substances with the release of a large amount of heat. Phosphorus combines most easily with oxygen, then with halogens, sulfur, and with many metals, and in the latter case, phosphides similar to nitrides are formed, for example, etc. All these properties are especially pronounced in white phosphorus; red phosphorus reacts less vigorously, black generally hardly enters into chemical interactions.