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Phosphorus exhibits allotropy (i.e., it has multiple forms in the same physical state); the physical constants given above are for the common white phosphorus. White phosphorus is an extremely poisonous, yellow to white, waxy, solid substance, nearly insoluble in water but very soluble in carbon disulfide. When exposed to air it ignites spontaneously, burning to form white fumes of phosphorus pentoxide, P2O5. Because of its toxicity and pyrophoric nature, phosphorus is stored underwater. Contact with the skin may cause burns. White phosphorus is phosphorescent (i.e., glows without emitting heat).
When white phosphorus is heated to about 250℃ in the absence of air, it changes into the more stable red phosphorus. This form appears as dull, reddish-brown cubic crystals or amorphous powder. Its specific gravity is 2.34. The red form is less dangerous than the white form, but should be handled with caution. It is insoluble in carbon disulfide and most other solvents. It does not ignite unless heated to about 200℃, does not phosphoresce, and is not poisonous. Another form of phosphorus is black phosphorus, a crystalline electrically conductive material similar to graphite in appearance. It was first prepared by P. W. Bridgman by heating white phosphorus to 200℃ under a pressure of 12,000 atmospheres. Its specific gravity is 2.70.
Natural Occurrence and Commercial Preparation
Biological Importance and Applications
Phosphorus is present in plants and animals. There is over 1 lb (454 grams) of phosphorus in the human body. It is a component of adenosine triphosphate (ATP), a fundamental energy source in living things. It is found in complex organic compounds in the blood, muscles, and nerves, and in calcium phosphate, the principal material in bones and teeth. Phosphorus compounds are essential in the diet. Organic phosphates, ferric phosphate, and tricalcium phosphate are added to foods. Dicalcium phosphate is added to animal feeds.
White phosphorus is used as a deoxidizing agent in the preparation of steel and phosphor bronze. It is also used in rat poisons, to make smoke screens and provide illumination (by burning) for warfare, and as an incendiary. Red phosphorus is used in making matches. The major use of phosphorus compounds is in fertilizers, especially in a mixture called superphosphate, obtained from phosphate minerals by sulfuric acid treatment; and in nitrophosphates. Phosphorus compounds are also used commercially in detergents, water softeners, pharmaceuticals, dentifrices, and in many other less important uses. Toxic nerve gases such as sarin contain phosphorus.
Phosphoric acid is primarily used in the production of phosphate compounds. It is also used in pickling metals, in sugar refining, and in soft drinks. Phosphorus forms a number of compounds with the halogens, e.g., the trichloride, PCl3, and the pentachloride, PCl5, both used as reagents. It also forms an oxychloride, POCl3. It reacts with sulfur to form a pentasulfide, P2S5, and a thiochloride, PSCl3, used in insecticides and oil additives. Phosphine, PH3, is a poisonous gas. Besides the pentoxide, phosphorus forms several other oxides; there are several acids other than the orthophosphoric acid noted above. Phosphorus also combines with various other nonmetals and with some metals.
P, a chemical element of group V of Mendeleev’s periodic system; a nonmetal. Atomic number, 15; atomic weight, 30.97376. Phosphorus occurs naturally only as the stable isotope 31P; six artificial radioisotopes have been obtained, namely 28P (half-life, 6.27 sec), 29P (4.45 sec), 30P (2.55 min), 32P (14.22 days), 33P (25 days), and 34P (12.5 sec). The isotope 32P is the most valuable; the energy associated with beta radiation is significant, and the isotope is used in both chemical and biological research as a tracer.
History. According to some sources, a method of preparing phosphorus was known as early as the 12th century by Arab alchemists. However, the generally accepted date for the discovery of the element is 1669, when H. Brand (Germany), upon roasting sand and a residue obtained from the evaporation of urine and then distilling the preparation in the absence of air, obtained a substance that glowed in the dark. Phosphorus was originally named cold fire; its present name comes from the Greek word phosphoros (light-bearing). The method of obtaining phosphorus soon became known to the German chemists J. Krafft and J. Kunckel and was published in 1682. In 1743, A. S. Marggraf developed a method for preparing phosphorus. Here, a mixture of lead chloride and urine was subjected to evaporation until dry and then heated until all the volatile products had been given off. The residue was mixed with charcoal in powder form and was subjected to distillation in a clay retort, and the phosphorus vapors were condensed in a container with water. The simplest method of preparing phosphorus, involving the roasting of bone ash with coal, was proposed in 1771 by C. Scheele. Phosphorus was established as an element by A. Lavoisier. The second half of the 19th century saw commercial production of phosphorus from phosphate rocks in retorts heated by coal-fired furnaces, which in the early 20th century gave way to electric furnaces.
Distribution. The average phosphorus content in the earth’s crust (clarke) is 9.3 × 10–2 percent by weight; it is 1.6 × 10–1 percent in intermediate rocks, 1.4 × 10–1 in basic rocks, 7 × 10–2 in granites and other acidic igneous rocks, and still lower—1.7 × 10–2—in ultrabasic rocks (mantle). In sedimentary rocks, the phosphorus content varies from 1.7 × 10–2 percent (sandstones) to 4 × 10–2 percent (carbonate rocks). Phosphorus figures in magmatic processes and migrates widely in the biosphere. The magmatic processes and migration are both associated with large accumulations of phosphorus, which form commercial deposits of apatites and phosphate rocks. Phosphorus is an exceptionally important biogenic element; it is accumulated by many organisms. Many processes of phosphorus concentration in the earth’s crust are associated with biogenic migration. Since phosphorus is taken in by living matter and is readily precipitated from water in the form of insoluble minerals, the concentration of the element in seawater is only 7 × 10–6 percent. Approximately 180 minerals of phosphorus are known, mainly phosphates of various types. Calcium phosphates are the most widely distributed.
Physical properties. Elemental phosphorus occurs in several allotropes, the main ones being the white, the red, and the black.
White phosphorus is a waxlike transparent substance with a characteristic odor; it is formed on condensation of phosphorus vapor. In the presence of trace amounts of red phosphorus, arsenic, iron, and other elements white phosphorus becomes yellow. For this reason, commercial white phosphorus is referred to as yellow phosphorus. White phosphorus exists in an α form and β form. The α modification occurs as crystals of the cubic system (a = 18.5 angstroms [Å]); it has a density of 1.828 g/cm3, a melting point of 44.1 °C, a boiling point of 280.5°C, and a heat of fusion of 2.5 kilojoules (kJ) per mole of P4 (0.6 kilocalorie [kcal] per mole of P4). The heat of vaporization is 58.6 kJ per mole of P 4 (14.0 kcal per mole of P4), and the vapor pressure at 25°C is 5.7 newtons (N) per sq m (0.043 mm of mercury). The coefficient of linear expansion in the temperature range 0°–44°C is 12.4 × 10–4, and the thermal conductivity at 25°C is 0.56 watt (W) per m·°K [1.1346 × 10–3cal/(cm · sec · °C)].
The electrical properties of α-white phosphorus are similar to those of dielectrics. The forbidden band comprises approximately 2.1 electron volts (eV), and the resistivity is 1.54 × 1011 ohm·cm. The α form is diamagnetic, with a magnetic susceptibility of –0.86 × 10–6. The Brinell hardness is 6 meganewtons (MN) per sq m (0.6 kilogram-force [kgf] per sq mm). The a form of white phosphorus is readily soluble in carbon disulfide but only sparingly soluble in liquid ammonia, benzene, and carbon tetra-chloride. At a temperature of –76.9°C and a pressure of 0.1 MN/m2 (1 kgf/cm2), the α form is converted into the β form (density, 1.88 g/cm3). If the pressure is increased to 1,200 MN/m2 (12,000 kgf/cm2), the transition will occur at 64.5°C. The β form occurs as crystals exhibiting double refraction; the structure of the crystals, however, has not been conclusively established. White phosphorus is poisonous. Since it ignites spontaneously in the air at a temperature of approximately 40°C, the element must be stored in water (the solubility in water at 25°C being 3.3 × 10–4 percent). If white phosphorus is heated in the absence of air at 250°–300°C for several hours, red phosphorus is obtained. The transition is exothermic and is accelerated by ultraviolet rays as well as by impurities (iodine, sodium, selenium). Standard commercial red phosphorus is nearly completely amorphous, its color ranging from dark brown to violet. With prolonged heating, it can be irreversibly converted to one of the crystal forms (triclinic, cubic), whose properties differ. The density varies from 2.0 to 2.4 g/cm3, the melting point from 585° to 610°C (at a pressure of tens of atmospheres), the sublimation point from 416° to 423°C, and the resistivity from 109 to 1014 ohm-cm. Red phosphorus does not ignite spontaneously in the air at temperatures up to 240°–250°C but does burn with friction or impact. It is insoluble in water, as well as in benzene and carbon disulfide, and soluble in phosphorus tribromide. At its sublimation point, red phosphorus is converted into vapor, which on cooling yields mainly white phosphorus.
Black phosphorus is formed by heating white phosphorus to 200°–220°C at a pressure of (1.2–1.7) × 103 MN/m2 [(12–17) × 103 kgf/cm2]. This conversion can be carried out without pressure provided that mercury and a small quantity of black phosphorus crystals (seeds) are present and that a temperature of 370°C is maintained for eight days. Black phosphorus occurs as crystals of orthorhombic structure (a= 3.31 Å, b= 4.38 Å, c= 10.50 Å); the lattice is constructed of corrugated layers with a pyramidal arrangement of atoms that is typical of phosphorus. The density is 2.69g/cm3, and the melting point approximately 1000°C at a pressure of 1.8 × 103 MN/m2 (18 × 103 kgf/cm2). Black phosphorus is outwardly similar to graphite. A semiconductor, it has a forbidden band of 0.33 eV at 25°C, a resistivity of 1.5 ohm-cm, and a temperature coefficient of resistivity of 0.0077. It is diamagnetic, with a specific magnetic susceptibility of –0.27 × 10–6. When heated to 560°–580°C under its own vapor pressure, black phosphorus is converted into red phosphorus. Black phosphorus is only slightly reactive; it does not catch fire easily when ignited and therefore can be subjected to mechanical working and exposure to the air.
Phosphorus has an atomic radius of 1.34 Å and ionic radii of 0.35 Å (P5+), 0.44 Å (P3+), and 1.86 Å (P3–).
Phosphorus atoms unite to form diatomic (P2), tetratomic (P4), and polymer molecules. Under normal conditions, polymer molecules containing long chains of interconnected P4 tetrahedra are the most stable. In the liquid form, in the solid form (white phosphorus), and in the vapor form (at temperatures below 800°C), phosphorus is made up of P4 molecules. Above 800°C, the P4 molecules dissociate into P2 molecules, which in turn are broken down into single atoms at temperatures above 2000°C. Only white phosphorus is composed of P4 molecules; all the other modifications are polymers.
Chemical properties. The electron configuration of the outer orbitals of the phosphorus atom is 3s23p3; oxidation states of +5, + 3, and –3 are most characteristic of phosphorus in its compounds. Like nitrogen, phosphorus forms mainly covalent bonds in its compounds. Compounds with ionic bonds, such as the phosphides Na3P and Ca3P2, are very few. Unlike nitrogen, phosphorus has free 3d orbitals with relatively low energies, which makes an increase in the coordination number possible and leads to the formation of donor-acceptor bonds.
Phosphorus is chemically active; white phosphorus exhibits the most activity, whereas red and black phosphorus are considerably more passive in chemical reactions. The oxidation of white phosphorus occurs through a chain reaction. The oxidation of phosphorus is usually accompanied by chemiluminescence. The burning of phosphorus in an excess of oxygen yields the pentoxide P4O10 (or P2O5); an insufficiency of oxygen results in the formation of mainly the trioxide P4O6 (or P2O3). The existence of P4O7, P4O8, P2O6, PO, and other phosphorus oxides in the vapors has been spectroscopically established. Phosphorus pentoxide is produced commercially by burning elemental phosphorus in an excess of dry air. Subsequent hydration of P4O10 yields orthophosphoric acid (H3PO4) and polyphosphoric acids (Hn+2PnO3n+1). In addition, phosphorus forms phosphorous acid (H3PO3), hypophosphoric acid (H4P2O6), and hypophosphorous acid (H3PO2), as well as two peracids, namely perphosphoric acid (H4P2O8) and monoperphosphoric acid (H3PO5). The salts of phosphoric acids—phosphates—have found wide application, whereas phosphites and hypophosphites are less widely used.
Phosphorus combines directly with all halogens, liberating a large amount of heat and forming trihalides (PX3, X being a halogen), pentahalides with the general formula PX5, and oxyhalides, for example, POX3. The fusion of phosphorus and sulfur at temperatures below 100°C yields solid solutions based on the two elements; temperatures above 100°C bring about the exothermic reaction for the formation of the crystalline sulfides P4S3, P4S5, P4S7, and P4S10. Of these, only P4S5 decomposes into P4S3 and P4S7 when heated above 200°C; the others melt without decomposition. The known oxysulfides of phosphorus are P2O3S2, P2O2S3, P4O4S3, P6O10S5, and P4O4S3. Compared with nitrogen, phosphorus is less capable of forming compounds with hydrogen. Hydrogen phosphide, or phosphine (PH3), and diphosphine (P2H4) can be obtained only by indirect means. The known compounds with nitrogen include the nitrides PN, P2N3, and P3N5—solid, chemically stable substances obtained by passing phosphorus vapor and nitrogen through an electric arc. Polymeric phosphonitrilic halides [(PNX2)], for example, polyphosphonitrilic chloride, are obtained by the reaction of halides of the general formula PX5 with ammonia under various conditions. Amidoimidophosphates as a rule are polymeric compounds that contain P—NH—P bonds in addition to P—O—P bonds.
At temperatures above 2000°C, phosphorus reacts with carbon to form the carbide PC3, a substance that is not soluble in ordinary solvents and that reacts with neither acids nor alkalies. When heated with metals, phosphorus forms phosphides.
Phosphorus forms numerous organophosphorus compounds.
Derivation. Elemental phosphorus is produced through the electrothermal reduction of naturally occurring phosphates (apatites or phosphate rocks) by coke at 1400°–1600°C in the presence of silica (quartz sand):
2Ca3(PO4)2 + 10C + nSiO2 = P4 + 10CO + 6CaO·nSiO2
Previously ground and dressed phosphorus-containing ore is mixed in given proportions with silica and coke and then charged into an electric furnace. Silica is necessary for reducing the temperature of the reaction and increasing the reaction rate. The latter effect is due to the bonding of the calcium oxide given off during the reduction process to form calcium silicate, which is continuously removed in the form of molten slag. Silicates and oxides of aluminum, magnesium, iron, and other impurities also enter the slag. Another component of the slag is ferrophosphorus (Fe2P, FeP, Fe3P), which is formed from the reaction of part of the reduced ifon with phosphorus. Ferrophosphorus, as well as the small amounts of the phosphides of manganese and other metals that are dissolved in it, is removed from the furnace as it accumulates for use in the manufacture of special steels.
Phosphorus vapor is discharged from the electric furnace along with gaseous by-products and volatile impurities (CO, SiF4, PH3, steam, products of the pyrolysis of organic impurities in the charge) at a temperature of 250°–350°C. After being purified of dust, the phosphorus-containing gases are fed into condensation units, where at a temperature not below 50°C liquid commercial white phosphorus is collected under water.
Methods of phosphorus production involving gaseous reducing agents and plasma reactors are being developed. These methods are designed to increase the yield through an increase of temperature to 2500°–3000°C, that is, to a point above the dissociation temperatures of naturally occurring phosphates and reducing gases, such as methane, that are used as transport gases in low-temperature plasma.
Use. Most of the phosphorus that is produced commercially is processed into phosphoric acid and into the phosphorus fertilizers and commercial salts (phosphates) based on the acid.
White phosphorus is used in incendiary bombs and devices and smoke bombs; red phosphorus is used in the production of matches. Phosphorus is used as a deoxidizing agent in the production of nonferrous metal alloys. The addition of phosphorus (up to 1 percent) increases the high-temperature strength of such alloys as fechral and Chromal. The element is also added to certain bronzes since it increases flowability and resistance to wear. The phosphides of metals, as well as of certain nonmetals (B, Si, As), are used in the production and doping of semiconductor materials. Phosphorus is used in preparing the chlorides and sulfides that serve as source material for the manufacture of phosphorus-containing plasticizers (tricresyl phosphate, tributyl phosphate), medications, and organophosphorus pesticides. The element is also used in preparing the chlorides and sulfides that are additives in lubricants and fuels.
Safety measures. White phosphorus and its compounds are highly toxic. Work with phosphorus requires that all equipment and apparatus be hermetically sealed. White phosphorus must be stored under water or in hermetically sealed metal containers. Safety regulations should be strictly observed when working with phosphorus.
L. V. KUBASOVA
Biological functions. Phosphorus is one of the most important biogenic elements; it is required for the vital activities of all organisms. The element is present in living cells in the form of orthophosphoric and pyrophosphoric acids and their derivatives, and it is also a component of nucleotides, nucleic acids, phosphoproteins, phospholipids, phosphorus esters of hydrocarbons, and many coenzymes and other organic compounds. Owing to the characteristic features of their chemical structure, phosphorus atoms, like sulfur atoms, are able to form energy-rich bonds in such high-energy compounds as adenosinetriphosphoric acid and creatine phosphate. During evolution, it was the phosphorus compounds that became the primary, universal guardians of genetic information and the energy carriers in all living systems. The importance of phosphorus compounds in organisms is also evident in the enzyme-catalyzed addition of the phosphoryl residue
to various organic compounds (phosphorylation); this addition is a prerequisite for the compounds’ participation in metabolism. Conversely, detachment of the phosphoryl residue (dephosphorylation) precludes metabolic activity. The enzymes of phosphorus metabolism comprise the kinases, phosphorylases, and phosphatases. The liver figures prominently in the conversion of phosphorus compounds in humans and animals. The metabolism of phosphorus compounds is regulated by hormones and vitamin D.
The phosphorus content (mg per 100 g dry matter) in the tissues of plants is 230–350; it is 400–1,800 in marine animals, 1,700–4,400 in terrestrial animals, and approximately 3,000 in bacteria. In humans, there are large amounts of phosphorus in bone tissue (slightly more than 5,000 mg per 100 g dry matter), brain tissue (approximately 4,000), and muscles (220–270). The daily human requirement of phosphorus is 1–1.2 g (children requiring more than adults). The food products richest in phosphorus are cheese, meat, eggs, and legumes (peas, beans). The phosphorus balance in organisms depends on the overall state of metabolism. Disruption of phosphorus metabolism leads to serious biochemical changes, first and foremost in energy metabolism. Phosphorus deficiency in humans and animals induces osteoporosis and other bone diseases; in plants, it causes phosphoric hunger. Phosphorus is obtained by organisms from its inorganic compounds that are present in the soil and dissolved in water. The element is extracted from the soil in the form of soluble phosphates by plants. Animals usually take in a sufficient amount of phosphorus in their food. When an organism dies, phosphorus reenters the soil and bottomset beds, thereby taking part in the earth’s cycle of matter. The important role of phosphorus in the regulation of metabolic processes derives from the extreme sensitivity of many enzyme systems in living cells to the action of organophosphorus compounds. This condition is used in medicine in developing therapeutic preparations and in agriculture in creating phosphorus fertilizers and effective insecticides. Many phosphorus compounds are extremely toxic, and certain organophosphorus compounds are considered war gases (sarin, soman, tabun). The radioisotope 32P is widely used in biology and medicine as a tracer in studying all types of metabolism and energy exchange in organisms.
N. N. CHERNOV
Poisoning. Poisoning can occur when working with white phosphorus; it is also possible during the thermoelectric sublimation of phosphorus and its compounds and during the production and use of the compounds. Organophosphorus compounds are highly toxic, having the effect of anticholinesterase. The element enters the body through the respiratory organs, gastrointestinal tract, and skin. The symptoms of acute poisoning include a burning sensation in the mouth and stomach, headache, weakness, nausea, and vomiting. After two or three days, jaundice manifests itself, and pain is felt in the epigastric region and right hypochondriac region. Chronic poisoning is attended by inflammation of the mucosa in the upper respiratory tract, symptoms of toxic hepatitis, disruption of calcium metabolism (development of osteoporosis; fragility and, sometimes, necrosis of the bone tissue, usually the mandible), and disorders of the cardiovascular and nervous systems.
First aid in cases of acute poisoning, occurring most often through the mouth, includes pumping of the stomach and administration of cathartics and enemas. Solutions of glucose, calcium chloride, and other substances are administered intravenously. For skin burns, the affected area is treated with solutions of blue vitriol or soda. Eyes should be rinsed with a 2-percent solution of sodium bicarbonate. Safety measures are important in working with phosphorus, as are personal and oral hygiene and regular (every six months) medical examinations.
Pharmaceuticals containing phosphorus (adenosinetriphosphoric acid, Phytin, calcium glycerophosphate, phosphrene) affect mainly the processes of endogenous metabolism; they are used against diseases of the muscles and nervous system, tuberculosis, anorexia, and anemia. Radioisotopes of phosphorus are used as tracers in studying metabolism and diagnosing diseases; they are also used in radiotherapy against tumors.
A. A. KASPAROV
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