PHOSPHORUS, P (lat. Phosphorus * a. phosphorus; n. Phosphor; f. phosphore; i. fosforo), - chemical element of group V periodic table Mendeleev, atomic number 15, atomic mass 30.97376. Natural phosphorus is represented by one stable isotope 31 R. 6 artificial ones are known radioactive isotopes phosphorus 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 () obtained a substance that glowed in the dark, called “cold fire”. The existence of phosphorus as a chemical element was proven in the early 70s. 18th century French chemist A. Lavoisier.
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 yellow, therefore marketable white phosphorus called yellow. There are 2 modifications of white phosphorus a-P has a densely packed cubic lattice 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 ; By electrical properties white phosphorus is close to dielectrics. At a temperature of 77.8°C and a pressure of 0.1 MPa, a-P transforms into b-P (rhombic lattice, density 1880 kg/m 3). Heating white phosphorus without air access at 250-300°C for several hours leads to the formation of a red modification. Ordinary commercial red phosphorus is practically amorphous, but upon prolonged heating it can transform 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 temperature 431°C), red phosphorus turns into 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 at normal pressure (at t 370°C), using as a catalyst, and also not a large number of black phosphorus for seeding. Black phosphorus - crystalline substance with a rhombic lattice (a=0.331, b=0.438 and c=1.05 nm), density 2690 kg/m 3, melting temperature 1000 °C; By appearance similar to graphite; semiconductor, diamagnetic. When heated to a temperature of 560-580°C and saturated vapor pressure, it turns into red phosphorus.
Phosphorus atoms combine into diatomic (P 2) and tetraatomic (P 4) polymer molecules. The most stable molecules under normal conditions are those 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 a covalent bond. Phosphorus is a chemically active element. Its white modification is characterized by the greatest activity, which spontaneously ignites at a temperature of about 40°C, therefore it is stored under a layer of water. Red phosphorus ignites when struck or rubbed. Black phosphorus is inactive and difficult to ignite when ignited. Phosphorus oxidation is usually accompanied by chemiluminescence. When phosphorus burns in an excess of oxygen, P 2 O 5 is formed, and when there is a deficiency, mainly P 2 O 3 is formed. 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, forming carbide (PC 3); when phosphorus is heated with metals - phosphides. White phosphorus and its compounds are highly toxic, MPC 0.03 mg/m3.
The average phosphorus content in the earth's crust (clarke) is 9.3.10 -2%, in ultrabasic rocks it is 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, forming 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 earth's crust. Over 180 minerals containing phosphorus are known.
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 phosphorus produced is processed into phosphoric acid and phosphorus fertilizers and technical salts obtained on its basis. Salts of phosphoric acids - phosphates, and to a slightly lesser extent - phosphites and hypophosphites are widely used. White phosphorus is used in the manufacture of incendiary and smoke projectiles; red - in match production.
Phosphorus is a fairly common chemical element on our planet. Its name translates as “luminous” because in pure form it glows brightly in the dark. This element was discovered completely by accident, by the alchemist Henning Brand, when he was trying to extract gold from urine. Thus, phosphorus became the first element that alchemists were able to obtain through their experiments.
Characteristics of phosphorus
It is chemically very active, so in nature it can only be found in the form of minerals - compounds with other elements, of which there are 190 species. The most important compound is calcium phosphate. Now many varieties of apatites are known, the most common of which is fluorapatite. From Apatity various types sedimentary rocks- phosphorites.
For living organisms, phosphorus is very important, since it is part of both plant and animal protein in the form various connections.
In plants, this element is found mainly in seed proteins, and in animal organisms - in various proteins in the blood, milk, brain cells and large amounts of phosphorus are found in the form of calcium phosphate in the bones of vertebrates.
Phosphorus exists in three allotropic modifications: white phosphorus, red and black. Let's take a closer look at them.
White phosphorus can be obtained by quickly cooling its vapor. Then a solid crystalline substance is formed, which in its pure form is absolutely colorless and transparent. White phosphorus sold for sale is usually slightly yellowish in color and closely resembles wax in appearance. In the cold, this substance becomes brittle, and at temperatures above 15 degrees it becomes soft and can be easily cut with a knife.
White phosphorus does not dissolve in water, but it responds well to organic solvents. In air it oxidizes very quickly (starts to burn) and at the same time glows in the dark. Actually, ideas about a luminous substance and detective stories about it are associated specifically with white phosphorus. It is a strong poison that is lethal even in small doses.
Red phosphorus is solid dark red color, which in its properties is strikingly different from those described above. It oxidizes in air very slowly, does not glow in the dark, lights up only when heated, it cannot be dissolved in organic solvents, and it is not poisonous. With strong heating, in which there is no access to air, it, without melting, turns into steam, from which, when cooled, white phosphorus is obtained. When both elements burn, phosphorus oxide is formed, which proves the presence of the same element in their composition. In other words, they are formed by one element - phosphorus - and are its allotropic modifications.
Black phosphorus is obtained from white phosphorus at 200 degrees Celsius under high pressure. It has a layered structure, a metallic luster and is similar in appearance to graphite. Of all hard species of this substance it is the least active.
Phosphorus compounds:
Phosphorus (the bringer of light) was first obtained by the Arab alchemist Ahad Behil in the 12th century. From European scientists first Phosphorus was discovered by the German Hennig Brant in 1669, while conducting experiments with human urine in attempts 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” (October 14, 1680; the work was published in 1693). Lavoisier later proved that phosphorus is a simple substance.
The phosphorus content in the earth's crust is 0.08% by weight - 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 apatite Ca 5 (PO 4) 3 (OH) and phosphorite Ca 3 (PO 4) 2.
Phosphorus plays an important role in the life of animals, plants and humans - it is part of such biological compounds as phospholipids, and is also present in proteins 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 an electronic configuration of the outer valence level similar to nitrogen (3s 2 3p 3), but phosphorus has less pronounced nonmetallic properties compared to nitrogen, which is explained by the presence of a free d-orbital, a larger atomic radius and lower ionization energy .
Reacting with others chemical elements, the phosphorus atom can exhibit an oxidation state from +5 to -3 (the most typical oxidation state is +5, the rest are quite rare).
In the ground (unexcited) state of the phosphorus atom on the outer energy level there are two paired electrons in the s-sublevel + 3 unpaired electrons in p-orbitals (the d-orbital is free). In the excited state, one electron moves from the s-sublevel to the d-orbital, which expands the valence capabilities of the phosphorus atom.
Rice. Transition of the phosphorus atom to an excited state.
Two phosphorus atoms combine to form a P2 molecule at a temperature of about 1000°C.
With more low temperatures phosphorus exists in tetraatomic P4 molecules as well as in more stable polymer P∞ molecules.
Allotropic modifications of phosphorus:
Of all the allotropic modifications of phosphorus, the most active is white phosphorus (P 4). Often in the equation chemical reactions they simply write P, not P4. 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 its properties in reactions with metals that occur when heated to form phosphides:
3Mg + 2P = Mg 3 P 2.
Phosphorus is reducing agent in reactions:
Phosphorus acts as both an oxidizing agent and a reducing agent in reactions disproportionation with aqueous solutions of alkalis when heated, forming (except for phosphine) hypophosphites (salts of hypophosphorous acid), in which it exhibits an uncharacteristic oxidation state of +1:
4P 0 +3KOH+3H 2 O = P -3 H 3 +3KH 2 P +1 O 2
YOU MUST REMEMBER: phosphorus does not react with other acids, except for the reactions indicated above.
Phosphorus is produced industrially by reducing it with coke from phosphorites (fluorapatates), which include calcium phosphate, by calcining them 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 high temperature silicon(IV) oxide displaces phosphorus(V) oxide from phosphate:
Ca 3 (PO 4) 2 + 3SiO 2 = 3CaSiO 3 + P 2 O 5.
Phosphorus (V) oxide is then reduced by coal to free phosphorus:
P 2 O 5 +5C = 2P+5CO.
Application of phosphorus:
The first mention of phosphorus ammunition dates back to the beginning of the 20th century - in 1916, grenades stuffed with white phosphorus appeared in England. During World War II, white phosphorus began to be used as one of the substances in the filling of incendiary bombs. IN last years Phosphorus weapons were actively used only by the American army, in particular in Iraq during the bombing of Fallujah.
Currently, phosphorus ammunition is understood as a type of incendiary or smoke ammunition filled with white phosphorus. There are several types of such weapons and ammunition, including aerial bombs, artillery shells, rockets (missiles), mortar shells, and hand grenades.
Unpurified white phosphorus is commonly called "yellow phosphorus". It is a flammable crystalline substance from light yellow to dark brown in color, which does not dissolve in water, and in air easily oxidizes and spontaneously ignites. White phosphorus as a chemical compound is very toxic (causes damage to bones, bone marrow, necrosis of the jaws).
A phosphorus bomb spreads a flammable substance whose combustion temperature exceeds 1200 °C. It burns with a dazzling, bright green flame and emits thick white smoke. Its distribution area can reach several hundred square meters. The combustion of the substance continues until the access of oxygen stops or all the phosphorus burns out.
To extinguish phosphorus, use water in large quantities(to reduce the temperature of the fire and convert phosphorus into a solid state) or a solution of copper sulfate ( copper sulfate), and after extinguishing the phosphorus is covered with wet sand. To protect against spontaneous combustion, yellow phosphorus is stored and transported under a layer of water (calcium chloride solution).
The use of white phosphorus gives a complex effect - not only serious physical injuries and slow death, but also psychological shock. The lethal dose of white phosphorus for an adult is 0.05-0.1 g. According to researchers, characteristic feature The use of this weapon results in charring of organic tissues, and when inhaling a burning mixture, burning out of the lungs.
Treatment of wounds caused by such weapons requires appropriately trained medical personnel. Special literature notes that inexperienced and untrained doctors can also receive phosphorus wounds when working with affected personnel.
Military use of ammunition containing white phosphorus against targets located in or near cities and other settlements, is prohibited under international agreements (Protocol III to the Convention on Certain Conventional Weapons).
From the history of the use of phosphorus bombs:
1916 In England, incendiary grenades filled with white phosphorus were supplied to arm troops.
The Second World War. White phosphorus began to be used as one of the substances in the filling of incendiary bombs.
In 1972, according to the conclusion of a special UN commission, incendiary weapons were conditionally classified as weapons of mass destruction.
1980 According to the Convention on Prohibitions or Restrictions on the Use specific types conventional weapons that may be considered to cause excessive injury or have an indiscriminate effect”, adopted by the UN, prohibits the use of incendiary weapons against civilian populations, and also prohibits the use of air-delivered incendiary weapons against military objectives in areas where civilian populations are concentrated.
In the 1980s, the Vietnamese People's Army used white phosphorus against Khmer Rouge guerrillas during the occupation of Kampuchea.
1982 155-mm artillery shells filled with white phosphorus were used by the Israeli army during the Lebanon War (in particular, during the siege of Beirut).
April 1984. In the area of the port of Bluefields, two Nicaraguan Contra saboteurs were blown up while trying to plant mines filled with white phosphorus.
June 1985. "Contra" passenger ship "Bluefields Express" and burned the ship with American phosphorus grenades.
1992 During the siege of Sarajevo, phosphorus shells were used by Bosnian Serb artillery.
2004 The Americans dropped bombs filled with this substance on Fallujah (Iraq).
In 2006, during the Second Lebanon War, artillery shells containing white phosphorus were used by the Israeli army.
year 2009. During Operation Cast Lead in the Gaza Strip, the Israeli army used smoke munitions containing white phosphorus.
year 2014. Semyonovka. The command of the anti-terrorist operation is committing war crimes against the civilian population of south-eastern Ukraine.
IN dark room or on the street at night, try this simple experiment. Not too hard, so that the match does not light up, strike it on the matchbox. You will notice that a glowing path from the match will be visible on the grater for a while. This glows white phosphorus. But everyone who remembers chemistry lessons high school, may say: “Excuse me, red, not white, phosphorus is used in the production of matches.” Right! There is no white phosphorus in the matchbox grater; there is red phosphorus, which, as a result of the reaction occurring between the red phosphorus located on the surface of the matchbox and the berthollet salt contained in the match head, heats up at the moment of friction and turns into white in a small amount.
Phosphorus can exist in several forms, or, as they say, in several modifications.
White phosphorus is a solid crystalline substance, and in its chemically pure form, white phosphorus crystals are completely colorless, transparent and refract light very well. In the light they quickly turn yellow and lose their transparency. Therefore, under normal conditions, phosphorus is very similar in appearance to wax, but is heavier (density of white phosphorus is 1.84). Phosphorus is fragile in the cold, but when room temperature relatively soft and easy to cut with a knife. At 44°C white phosphorus melts, and at 280.5°C it boils. White phosphorus, oxidized by oxygen in the air, glows in the dark and easily ignites when slightly heated, for example from friction.
The ignition temperature of completely dry and pure phosphorus is close to the temperature human body. Therefore, it is stored only under water. First world war white phosphorus was used as an incendiary material in artillery shells, aerial bombs, grenades, and bullets.
Red phosphorus, in contrast to white, or yellow, as it is sometimes called, is not poisonous, does not oxidize in air, does not glow in the dark, does not dissolve in carbon disulfide and ignites only at 260 ° C. Red phosphorus is obtained from white phosphorus by prolonged heating without air access at 250-300°C.
Joseph Wright's painting "The Alchemist Discovering Phosphorus" supposedly describes Hennig Brand's discovery of phosphorus
In search of the elixir of youth and attempts to obtain gold, the 17th century alchemist Genning Brand from Hamburg tried to make a “philosopher’s stone” from urine. For this purpose, he evaporated a large amount of it and the syrupy residue obtained after evaporation was subjected to strong calcination in a mixture with sand and charcoal without air access.
As a result, Brand received a substance with extraordinary properties: it glowed in the dark; thrown into boiling water, it released vapors that ignited in the air, releasing thick white smoke that dissolved in the water to form acid.
There was enormous interest in the new substance, and Brand hoped to make a hefty profit from his discovery: it was not for nothing that he was a former Hamburg merchant. Keeping the manufacturing method in the strictest confidence, Brand showed the new substance for money and sold it to those who wanted it. in small portions only for pure gold. After some time, Brand also sold the secret of making phosphorus to the Dresden chemist Kraft, who, like Brand, began to travel around the palaces of influential people, showing phosphorus for money, making a huge fortune.
After the discovery of phosphorus, its ability to glow in the dark was again used, but for different purposes. This time, representatives of religious cults began to trade in phosphorus. Recipes for using phosphorus were very diverse. For example, wax or paraffin was added to melted but already thickened a small amount of white phosphorus. The resulting mixture was used to mold pencils, which were used to write on the walls of churches and icons. At night, “mysterious inscriptions” were visible. Phosphorus, slowly oxidizing, glowed, and paraffin, protecting it from rapid oxidation, increased the duration of the phenomenon.
White phosphorus was dissolved in benzene or carbon disulfide. The resulting solution was used to moisten the wicks of candles or lamps. After the solvent evaporated, the white phosphorus ignited, and the wick ignited from it. This is how a “miracle” called “self-ignition of candles” was fabricated.
One of interesting connections Phosphorus is a gaseous phosphine, the peculiarity of which is that it is highly flammable in air. This property of phosphine explains the appearance of swamp, will-o'-the-wisp, or grave-lights. There really are fires in swamps and fresh graves. This is not fantasy or fiction. On warm, dark nights, pale bluish, faintly flickering lights are sometimes observed on fresh graves. It is the phosphine that “burns.” Phosphine is formed during the decay of dead plant and animal organisms.
