Phosphorus ( , ) is the chemical element that has the symbol P and atomic number 15. A multivalent nonmetal of the nitrogen group, phosphorus is commonly found in inorganic phosphate rocks. Elemental phosphorus exists in two major forms - white phosphorus and red phosphorus. Although the term "phosphorescence", meaning glow after illumination, derives from phosphorus, glow of phosphorus originates from oxidation of the white (but not red) phosphorus and should be called chemiluminescence.

Due to its high reactivity, phosphorus is never found as a free element in nature on Earth. The first form of phosphorus to be discovered (white phosphorus, in 1669) emits a faint glow upon exposure to oxygen — hence its name given from Greek mythology, meaning "light-bearer" (Latin Lucifer), referring to the "Morning Star", the planet Venus.

Phosphorus is a component of DNA, RNA, ATP, and also the phospholipids which form all cell membranes. It is thus an essential element for all living cells. The most important commercial use of phosphorus-based chemicals is the production of fertilizers.

Phosphorus compounds are also widely used in explosives, nerve agents, friction matches, fireworks, pesticides, toothpaste and detergents.

Physical properties

Glow from white phosphorus

In 1669, German alchemist Hennig Brandt attempted to distil some kind of "life essence" from his urine, and in the process produced a white material that glowed in the dark. The phosphorus had in fact been produced from inorganic phosphate, which is a significant component of dissolved urine solids. White phosphorus is highly reactive and gives off a faint greenish glow upon uniting with oxygen. The glow observed by Brand was actually caused by the very slow burning of the phosphorus, but as he neither saw flame nor felt any heat he did not recognize it as burning.

It was known from early times that the glow would persist for a time in a stoppered jar but then cease. Robert Boyle in the 1680s ascribed it to "debilitation" of the air; in fact, it is oxygen being consumed. By the 18th century, it was known that in pure oxygen, phosphorus does not glow at all; there is only a range of partial pressure at which it does. Heat can be applied to drive the reaction at higher pressures.

In 1974, the glow was explained by R. J. van Zee and A. U. Khan. A reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived molecules HPO and P₂O₂ that both emit visible light. The reaction is slow and only very little of the intermediates are required to produce the luminescence, hence the extended time the glow continues in a stoppered jar.

Although the term phosphorescence is derived from phosphorus, the reaction which gives phosphorus its glow is properly called chemiluminescence (glowing due to a cold chemical reaction), not phosphorescence (re-emitting light that previously fell onto a substance and excited it).

Phosphorescence is the slow decay of a metastable electronic state to a lower energy state through emission of light. The decay is slow because the transition from the excited to the lower state requires a spin flip, making it classically forbidden. Often it involves a transition from an excited triplet state to a singlet ground state. The metastable excited state may have been populated by thermal excitations or some light source. Since phosphorescence is slow, it persists for some time after the exciting source is removed. In contrast, chemiluminescence occurs when the product molecules of a chemical reaction (HPO and P₂O₂ in this case) leave the reaction in an electronically excited state. These excited molecules then release their excess energy in the form of light. The frequency (colour) of the light emitted is proportional to the energy difference of the two electronic states involved.

Allotropes

Phosphorus has several forms (allotropes) which have strikingly different properties. The two most common allotropes are white phosphorus and red phosphorus. Red phosphorus is an intermediate phase between white and violet phosphorus. Another form, scarlet phosphorus, is obtained by allowing a solution of white phosphorus in carbon disulfide to evaporate in sunlight. Black phosphorus is obtained by heating white phosphorus under high pressures (about 12,000 atmospheres). In appearance, properties, and structure, it resembles graphite, being black and flaky, a conductor of electricity, and has puckered sheets of linked atoms. Another allotrope is diphosphorus; it contains a phosphorus dimer as a structural unit and is highly reactive.

White phosphorus has two forms, low-temperature β form and high-temperature α form. They both contain a phosphorus tetrahedron as a structural unit, in which each atom is bound to the other three atoms by a single bond. This tetrahedron is also present in liquid and gaseous phosphorus up to the temperature of 800 °C when it starts decomposing to molecules. White phosphorus is the least stable, the most reactive, more volatile, less dense, and more toxic than the other allotropes. The toxicity of white phosphorus led to its discontinued use in matches. White phosphorus is thermodynamically unstable at normal condition and will gradually change to red phosphorus. This transformation, which is accelerated by light and heat, makes white phosphorus almost always contain some red phosphorus and therefore appear yellow. For this reason, it is also called yellow phosphorus. It glows greenish in the dark (when exposed to oxygen), is highly flammable and pyrophoric (self-igniting) upon contact with air as well as toxic (causing severe liver damage on ingestion). Because of pyrophoricity, white phosphorus is used as an additive in napalm. The odour of combustion of this form has a characteristic garlic smell, and samples are commonly coated with white "(di)phosphorus pentoxide", which consists of P₄O₁₀ tetrahedra with oxygen inserted between the phosphorus atoms and at their vertices. White phosphorus is insoluble in water but soluble in carbon disulfide.

The white allotrope can be produced using several different methods. In one process, calcium phosphate, which is derived from phosphate rock, is heated in an electric or fuel-fired furnace in the presence of carbon and silica. Elemental phosphorus is then liberated as a vapour and can be collected under phosphoric acid. This process is similar to the first synthesis of phosphorus from calcium phosphate in urine.

In the red phosphorus, one of the P₄ bonds is broken, and one additional bond is formed with a neighbouring tetrahedron resulting in a more chain-like structure. Red phosphorus may be formed by heating white phosphorus to 250 °C (482 °F) or by exposing white phosphorus to sunlight. Phosphorus after this treatment exists as an amorphous network of atoms which reduces strain and gives greater stability; further heating results in the red phosphorus becoming crystalline. Therefore red phosphorus is not a certain allotrope, but rather an intermediate phase between the white and violet phosphorus, and most of its properties have a range of values. Red phosphorus does not catch fire in air at temperatures below 260 °C, whereas white phosphorus ignites at about 30 °C.

Violet phosphorus is a thermodynamic stable form of phosphorus which can be produced by day-long temper of red phosphorus above 550 °C. In 1865, Hittorf discovered that when phosphorus was recrystallized from molten lead, a red/purple form is obtained. Therefore this form is sometimes known as "Hittorf's phosphorus" (or violet or α-metallic phosphorus).

Black phosphorus is the least reactive allotrope and the thermodynamic stable form below 550 °C. It is also known as β-metallic phosphorus and has a structure somewhat resembling that of graphite. High pressures are usually required to produce black phosphorus, but it can also be produced at ambient conditions using metal salts as catalysts.

The diphosphorus allotrope, P₂, is stable only at high temperatures. The dimeric unit contains a triple bond and is analogous to N₂. The diphosphorus allotrope (P₂) can be obtained normally only under extreme conditions (for example, from P₄ at 1100 kelvin). Nevertheless, some advancements were obtained in generating the diatomic molecule in homogeneous solution, under normal conditions with the use by some transitional metal complexes (based on, for example, tungsten and niobium).

Properties of some allotropes of phosphorus

Form	white(α)	white(β)	violet	black
Symmetry	Body-centred cubic	Triclinic	Monoclinic	Orthorhombic
Pearson symbol		aP24	mP84	oS8
Space group	I-43m	P-1 No.2	P2/c No.13	Cmca No.64
Density (g/cm3)	1.828	1.88	2.36	2.69
Bandgap (eV)	2.1		1.5	0.34
Refractive index	1.8244		2.6	2.4

Isotopes

Although twenty-three isotopes of phosphorus are known (all possibilities from ²⁴P up to ⁴⁶P), only ³¹P, with spin 1/2, is stable and is therefore present at 100% abundance. The half-integer spin and high abundance of ³¹P make it useful for nuclear magnetic resonance studies of biomolecules, particularly DNA.

Two radioactive isotopes of phosphorus have half-lives which make them useful for scientific experiments. ³²P has a half-life of 14.262 days and ³³P has a half-life of 25.34 days. Biomolecules can be "tagged" with a radioisotope to allow for the study of very dilute samples.

Radioactive isotopes of phosphorus include

• ³²P, a beta-emitter (1.71 MeV) with a half-life of 14.3 days which is used routinely in life-science laboratories, primarily to produce radiolabeled DNA and RNA probe, e.g. for use in Northern blots or Southern blots. Because the high energy beta particles produced penetrate skin and corneas, and because any ³²P ingested, inhaled, or absorbed is readily incorporated into bone and nucleic acids, Occupational Safety and Health Administration in the United States, and similar institutions in other developed countries require that a lab coat, disposable gloves, and safety glasses or goggles be worn when working with ³²P, and that working directly over an open container be avoided in order to protect the eyes. Monitoring personal, clothing, and surface contamination is also required. In addition, due to the high energy of the beta particles, shielding this radiation with the normally used dense materials (e.g. lead), gives rise to secondary emission of X-rays via a process known as Bremsstrahlung, meaning braking radiation. Therefore shielding must be accomplished with low density materials, e.g. Plexiglas, Lucite, plastic, wood, or water.
• ³³P, a beta-emitter (0.25 MeV) with a half-life of 25.4 days. It is used in life-science laboratories in applications in which lower energy beta emissions are advantageous such as DNA sequencing.

Chemical properties

• Hydrides: PH₃,P₂H₄
• Halides: PBr₅, PBr₃, PCl₃, PI₃
• Oxides:P₄O₆, P₄O₁₀
• Sulfides: P₄S₆, P₄S₁₀
• Acids: H₃PO₂, H₃PO₄
• Phosphates: ₃PO₄, Ca₃₂, FePO₄, Fe₃₂, Na₃PO₄, Ca₂, KH₂PO₄
• Phosphides: Ca₃P₂, GaP, Zn₃P₂ Cu₃P
• Organophosphorus and organophosphates: Lawesson's reagent, Parathion, Sarin, Soman, Tabun, Triphenyl phosphine, VX nerve gas

Chemical bonding

Because phosphorus is just below nitrogen in the periodic table, the two elements share many of their bonding characteristics. For instance, phosphine, PH₃, is an analogue of ammonia, NH₃. Phosphorus, like nitrogen, is trivalent in this molecule.

The "trivalent" or simple 3-bond view is the pre-quantum mechanical Lewis structure, which although somewhat of a simplification from a quantum chemical point of view, illustrates some of the distinguishing chemistry of the element. In quantum chemical valence bond theory, the valence electrons are seen to be in mixtures of four s and p atomic orbitals, so-called hybrids. In this view, the three unpaired electrons in the three 3p orbitals combine with the two electrons in the 3s orbital to form three electron pairs of opposite spin, available for the formation of three bonds. The remaining hybrid orbital contains two paired non-bonding electrons, which show as a lone pair in the Lewis structure.

The phosphorus cation is very similar to the nitrogen cation. In the same way that nitrogen forms the tetravalent ammonium ion, phosphorus can form the tetravalent phosphonium ion, and form salts such as phosphonium iodide [PH₄]⁺[I⁻].

Like other elements in the third or lower rows of the periodic table, phosphorus atoms can expand their valence to make penta- and hexavalent compounds. The phosphorus chloride molecule is an example. When the phosphorus ligands are not identical, the more electronegative ligands are located in the apical positions and the least electronegative ligands are located in the axial positions.

With strongly electronegative ions, in particular fluorine, hexavalency as in PF₆⁻ occurs as well. This octahedral ion is isoelectronic with SF₆. In the bonding the six octahedral sp³d² hybrid atomic orbitals play an important role.

Before extensive computer calculations were feasible, it was generally assumed that the nearby d orbitals in the n = 3 shell were the obvious cause of the difference in binding between nitrogen and phosphorus (i.e., phosphorus had 3d orbitals available for 3s and 3p shell bonding electron hybridisation, but nitrogen did not). However, in the early eighties the German theoretical chemist Werner Kutzelnigg found from an analysis of computer calculations that the difference in binding is more likely due to differences in character between the valence 2p and valence 3p orbitals of nitrogen and phosphorus, respectively. The 2s and 2p orbitals of first row atoms are localized in roughly the same region of space, while the 3p orbitals of phosphorus are much more extended in space. The violation of the octet rule observed in compounds of phosphorus is then due to the size of the phosphorus atom, and the corresponding reduction of steric hindrance between its ligands. In modern theoretical chemistry, Kutzelnigg's analysis is generally accepted.

The simple Lewis structure for the trigonal bipyramidal PCl₅ molecule contains five covalent bonds, implying a hypervalent molecule with ten valence electrons contrary to the octet rule.

An alternate description of the bonding, however, respects the octet rule by using 3-centre-4-electron bonds. In this model the octet on the P atom corresponds to six electrons which form three Lewis (2c-2e) bonds to the three equatorial Cl atoms, plus the two electrons in the 3-centre Cl-P-Cl bonding molecular orbital for the two axial Cl electrons. The two electrons in the corresponding nonbonding molecular orbital are not included because this orbital is localized on the two Cl atoms and does not contribute to the electron density on the phosphorus atom. (However, it should always be remembered that the octet rule is not some universal rule of chemical bonding, and while many compounds obey it, there are many elements to which it does not apply).

Phosphine, diphosphine and phosphonium salts

Phosphine (PH₃) and arsine (AsH₃) are structural analogues with ammonia (NH₃) and form pyramidal structures with the phosphorus or arsenic atom in the centre bound to three hydrogen atoms and one lone electron pair. Both are colourless, ill-smelling, toxic compounds. Phosphine is produced in a manner similar to the production of ammonia. Hydrolysis of calcium phosphide, Ca₃P₂, or calcium nitride, Ca₃N₂ produces phosphine or ammonia, respectively. Unlike ammonia, phosphine is unstable and it reacts instantly with air giving off phosphoric acid clouds. Arsine is even less stable. Although phosphine is less basic than ammonia, it can form some phosphonium salts (like PH₄I), analogues of ammonium salts, but these salts immediately decompose in water and do not yield phosphonium (PH₄⁺) ions. Diphosphine (P₂H₄ or H₂P-PH₂) is an analogue of hydrazine (N₂H₄) that is a colourless liquid which spontaneously ignites in air and can disproportionate into phosphine and complex hydrides.

Halides

The trihalides PF₃, PCl₃, PBr₃ and PI₃ and the pentahalides, PCl₅ and PBr₅ are all known and mixed halides can also be formed. The trihalides can be formed simply by mixing the appropriate stoichiometric amounts of phosphorus and a halogen.For safety reasons, however, PF₃ is typically made by reacting PCl₃ with AsF₅ and fractional distillation because the direct reaction of phosphorus with fluorine can be explosive. The pentahalides, PX₅, are synthesized by reacting excess halogen with either elemental phosphorus or with the corresponding trihalide. Mixed phosphorus halides are unstable and decompose to form simple halides. Thus 5PF₃Br₂ decomposes into 3PF₅ and 2PBr₅.

Oxides and oxyacids

Phosphorus oxide, P₄O₆ (also called tetraphosphorus hexoxide) and phosphorus oxide, P₄O₁₀ (or tetraphosphorus decoxide) are acid anhydrides of phosphorus oxyacids and hence readily react with water. P₄O₁₀ is a particularly good dehydrating agent that can even remove water from nitric acid, HNO₃. The structure of P₄O₆ is like that of P₄ with an oxygen atom inserted between each of the P-P bonds. The structure of P₄O₁₀ is like that of P₄O₆ with the addition of one oxygen bond to each phosphorus atom via a double bond and protruding away from the tetrahedral structure.

Phosphorous oxyacids can have acidic protons bound to oxygen atoms and nonacidic protons which are bonded directly to the phosphorus atom. Although many oxyacids of phosphorus are formed, only six are important (see table), and three of them, hypophosphorous acid, phosphorous acid and phosphoric acid are particularly important ones.

Oxidation state	Formula	Name	Acidic protons	Compounds
+1	H3PO2	hypophosphorous acid	1	acid, salts
+3	H3PO3	(ortho)phosphorous acid	2	acid, salts
+5	(HPO3)n	metaphosphoric acids	n	salts (n=3,4)
+5	H5P3O10	triphosphoric acid	3	salts
+5	H4P2O7	pyrophosphoric acid	4	acid, salts
+5	H3PO4	(ortho)phosphoric acid	3	acid, salts

Spelling and etymology

The name Phosphorus in Ancient Greece was the name for the planet Venus and is derived from the Greek words (φως = light, φορέω = carry) which roughly translates as light-bringer or light carrier. (In Greek mythology, Hesperus (evening star) and Eosphorus (dawnbearer) are close homologues, and also associated with Phosphorus-the-planet).

According to the Oxford English Dictionary, the correct spelling of the element is phosphorus. The word phosphorous is the adjectival form of the P³⁺ valence: so, just as sulfur forms sulfurous and sulfuric compounds, phosphorus forms phosphorous compounds (see, e.g., phosphorous acid) and P⁵⁺ valency phosphoric compounds (see, e.g., phosphoric acids and phosphates).

History and discovery

The discovery of phosphorus is credited to the German alchemist Hennig Brand in 1669, although other chemists might have discovered phosphorus around the same time. Brand experimented with urine, which contains considerable quantities of dissolved phosphates from normal metabolism. Working in Hamburg, Brand attempted to create the fabled philosopher's stone through the distillation of some salts by evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. His process originally involved letting urine stand for days until it gave off a terrible smell. Then he boiled it down to a paste, heated this paste to a high temperature, and led the vapours through water, where he hoped they would condense to gold. Instead, he obtained a white, waxy substance that glowed in the dark. Brand had discovered phosphorus, the first element discovered since antiquity. We now know that Brand produced ammonium sodium hydrogen phosphate, (NH₄)NaHPO₄. While the quantities were essentially correct (it took about 1,100 L of urine to make about 60 g of phosphorus), it was unnecessary to allow the urine to rot. Later scientists would discover that fresh urine yielded the same amount of phosphorus.

Since that time, phosphors and phosphorescence were used loosely to describe substances that shine in the dark without burning. However, as mentioned above, even though the term phosphorescence was originally coined as a term by analogy with the glow from oxidation of elemental phosphorus, is now reserved for another fundamentally different process—re-emission of light after illumination.

Phosphorus was gradually recognized as a chemical element in its own right at the emergence of the atomic theory that gradually occurred in the late part of the 18th century and the early 19th century (see John Dalton for more history).

Brand at first tried to keep the method secret, but later sold the recipe for 200 thaler to D Krafft from Dresden, who could now make it as well, and toured much of Europe with it, including England, where he met with Robert Boyle. The secret that it was made from urine leaked out and first Johann Kunckel (1630-1703) in Sweden (1678) and later Boyle in London (1680) also managed to make phosphorus. Boyle states that Krafft gave him no information as to the preparation of phosphorus other than that it was derived from "somewhat that belonged to the body of man". This gave Boyle a valuable clue, however, so that he, too, managed to make phosphorus, and published the method of its manufacture. Later he improved Brand's process by using sand in the reaction (still using urine as base material),

4 NaPO₃ + 2 SiO₂ + 10 C → 2 Na₂SiO₃ + 10 CO + P₄

Robert Boyle was the first to use phosphorus to ignite sulfur-tipped wooden splints, forerunners of our modern matches, in 1680.

In 1769 Johan Gottlieb Gahn and Carl Wilhelm Scheele showed that calcium phosphate (Ca₃(PO₄)₂) is found in bones, and they obtained phosphorus from bone ash. Antoine Lavoisier recognized phosphorus as an element in 1777. Bone ash was the major source of phosphorus until the 1840s. Phosphate rock, a mineral containing calcium phosphate, was first used in 1850 and following the introduction of the electric arc furnace in 1890, this became the only source of phosphorus. Phosphorus, phosphates and phosphoric acid are still obtained from phosphate rock. Phosphate rock is a major feedstock in the fertilizer industry.

Early matches used white phosphorus in their composition, which was dangerous due to its toxicity. Murders, suicides and accidental poisonings resulted from its use. (An apocryphal tale tells of a woman attempting to murder her husband with white phosphorus in his food, which was detected by the stew giving off luminous steam). In addition, exposure to the vapours gave match workers a severe necrosis of the bones of the jaw, the infamous "phossy jaw." When a safe process for manufacturing red phosphorus was discovered, with its far lower flammability and toxicity, laws were enacted, under the Berne Convention , requiring its adoption as a safer alternative for match manufacture.

Occurrence

Due to its reactivity with air and many other oxygen-containing substances, phosphorus is not found free in nature but it is widely distributed in many different minerals.

Phosphate rock, which is partially made of apatite (an impure tri-calcium phosphate mineral), is an important commercial source of this element. About 50 percent of the global phosphorus reserves are in the Arab nations. Large deposits of apatite are located in China, Russia, Morocco, Florida, Idaho, Tennessee, Utah, and elsewhere. Albright and Wilson in the United Kingdom and their Niagara Falls plant, for instance, were using phosphate rock in the 1890s and 1900s from Connetable, Tennessee and Florida; by 1950 they were using phosphate rock mainly from Tennessee and North Africa. In the early 1990s Albright and Wilson's purified wet phosphoric acid business was being adversely affected by phosphate rock sales by China and the entry of their long-standing Moroccan phosphate suppliers into the purified wet phosphoric acid business.

In 2007, at the current rate of consumption, the supply of phosphorus was estimated to run out in 345 years. However, scientists are now claiming that a "Peak Phosphorus" will occur in 30 years and that "At current rates, reserves will be depleted in the next 50 to 100 years."

Production

White phosphorus was first made commercially, for the match industry in the 19th century, by distilling off phosphorus vapour from precipitated phosphates, mixed with ground coal or charcoal, which was heated in an iron pot, in retort. The precipitated phosphates were made from ground-up bones that had been de-greased and treated with strong acids. Carbon monoxide and other flammable gases produced during the reduction process were burnt off in a flare stack.

This process became obsolete in the late 1890s when the electric arc furnace was adapted to reduce phosphate rock. Calcium phosphate (phosphate rock), mostly mined in Florida and North Africa, can be heated to 1,200-1,500 °C with sand, which is mostly SiO₂, and coke (impure carbon) to produce vaporized tetraphosphorus, P₄, (mp. 44.2 C) which is subsequently condensed into a white powder under water to prevent oxidation. Even under water, white phosphorus is slowly converted to the more stable red phosphorus allotrope (mp. 597 C). Both the white and red allotropes of phosphorus are insoluble in water.

The electric furnace method allowed production to increase to the point where phosphorus could be used in weapons of war. In World War I it was used in incendiaries, smoke screens and tracer bullets. A special incendiary bullet was developed to shoot at hydrogen-filled Zeppelins over Britain (hydrogen being highly inflammable if it can be ignited). During World War II, Molotov cocktails of benzene and phosphorus were distributed in Britain to specially selected civilians within the British resistance operation, for defence; and phosphorus incendiary bombs were used in war on a large scale. Burning phosphorus is difficult to extinguish and if it splashes onto human skin it has horrific effects (see precautions below).

Today phosphorus production is larger than ever. It is used as a precursor for various chemicals, in particular the herbicide glyphosate sold under the brand name Roundup. Production of white phosphorus takes place at large facilities and it is transported heated in liquid form. Some major accidents have occurred during transportation, train derailments at Brownston, Nebraska and Miamisburg, Ohio led to large fires. The worst accident in recent times was an environmental one in 1968 when phosphorus spilled into the sea from a plant at Placentia Bay, Newfoundland.

Applications

Widely used compounds	Use
Ca(H2PO4)2•H2O	Baking powder & fertilizers
CaHPO4•2H2O	Animal food additive, toothpowder
H3PO4	Manufacture of phosphate fertilizers
PCl3	Manufacture of POCl3 and pesticides
POCl3	Manufacturing plasticizer
P4S10	Manufacturing of additives and pesticides
Na5P3O10	Detergents

Phosphorus, being an essential plant nutrient, finds its major use as a constituent of fertilizers for agriculture and farm production in the form of concentrated phosphoric acids, which can consist of 70% to 75% P₂O₅. Global demand for fertilizers led to large increase in phosphate (PO₄^3-) production in the second half of the 20th century. Due to the essential nature of phosphorus to living organisms, the low solubility of natural phosphorus-containing compounds, and the slow natural cycle of phosphorus, the agricultural industry is heavily reliant on fertilizers which contain phosphate, mostly in the form of superphosphate of lime. Superphosphate of lime is a mixture of two phosphate salts, calcium dihydrogen phosphate Ca(H₂PO₄)₂ and calcium sulfate dihydrate CaSO₄•2H₂O produced by the reaction of sulfuric acid and water with calcium phosphate.

• Phosphorus is widely used to make organophosphorus compounds, through the intermediates phosphorus chlorides and two phosphorus sulfides: phosphorus pentasulfide, and phosphorus sesquisulfide. Organophosphorus compounds have many applications, including in plasticizers, flame retardants, pesticides, extraction agents, and water treatment.
• Phosphorus is also an important component in steel production, in the making of phosphor bronze, and in many other related products.
• Phosphates are utilized in the making of special glasses that are used for sodium lamps.
• Bone-ash, calcium phosphate, is used in the production of fine china.
• Sodium tripolyphosphate made from phosphoric acid is used in laundry detergents in some countries, but banned for this use in others.
• Phosphoric acid made from elemental phosphorus is used in food applications such as some soda beverages. The acid is also a starting point to make food grade phosphates. These include mono-calcium phosphate which is employed in baking powder and sodium tripolyphosphate and other sodium phosphates. Among other uses these are used to improve the characteristics of processed meat and cheese. Others are used in toothpaste. Trisodium phosphate is used in cleaning agents to soften water and for preventing pipe/boiler tube corrosion.
• White phosphorus, called "WP" (slang term "Willie Peter") is used in military applications as incendiary bomb, for smoke-screening as smoke pots and smoke bombs, and in tracer ammunition. It is also a part of an obsolete M34 White Phosphorus US hand grenade. This multipurpose grenade was mostly used for signalling, smoke screens and inflammation; it could also cause severe burns and had a psychological impact on the enemy.
• Red phosphorus is essential for manufacturing matchbook strikers, flares, safety matches, pharmaceutical grade and street methamphetamine, and is used in cap gun caps.
• Phosphorus sesquisulfide is used in heads of strike-anywhere matches.
• In trace amounts, phosphorus is used as a dopant for n-type semiconductors.
• ³²P and ³³P are used as radioactive tracers in biochemical laboratories (see Isotopes).

Biological role

Phosphorus is a key element in all known forms of life. Inorganic phosphorus in the form of the phosphate PO₄^3- plays a major role in biological molecules such as DNA and RNA where it forms part of the structural framework of these molecules. Living cells also use phosphate to transport cellular energy in the form of adenosine triphosphate (ATP). Nearly every cellular process that uses energy obtains it in the form of ATP. ATP is also important for phosphorylation, a key regulatory event in cells. Phospholipids are the main structural components of all cellular membranes. Calcium phosphate salts assist in stiffening bones.

Every cell has a membrane that separates it from its surrounding environment. Biological membranes are made from a phospholipid matrix and proteins, typically in the form of a bilayer. Phospholipids are derived from glycerol, such that two of the glycerol hydroxyl (OH) protons have been replaced with fatty acids as an ester, and the third hydroxyl proton has been replaced with phosphate bonded to another alcohol.

An average adult human contains about 0.7 kg of phosphorus, about 85-90% of which is present in bones and teeth in the form of apatite, and the remainder in soft tissues and extracellular fluids (~1%). The phosphorus content increases from about 0.5 weight% in infancy to 0.65-1.1 weight% in adults. Average phosphorus concentration in the blood is about 0.4 g/L, about 70% of that is organic and 30% inorganic phosphates. A well-fed adult in the industrialized world consumes and excretes about 1-3 g of phosphorus per day, with consumption in the form of inorganic phosphate and phosphorus-containing biomolecules such as nucleic acids and phospholipids; and excretion almost exclusively in the form of urine phosphate ion. Only about 0.1% of body phosphate circulates in the blood, but this amount reflects the amount of phosphate available to soft tissue cells.

In medicine, low phosphate syndromes are caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes which draw phosphate from the blood (such as re-feeding after malnutrition) or pass too much of it into the urine. All are characterized by hypophosphatemia (see article for medical details), which is a condition of low levels of soluble phosphate levels in the blood serum, and therefore inside cells. Symptoms of hypophosphatemia include muscle and neurological dysfunction, and disruption of muscle and blood cells due to lack of ATP. Too much phosphate can lead to diarrhoea and calcification (hardening) of organs and soft tissue, and can interfere with the body's ability to use iron, calcium, magnesium, and zinc.

Phosphorus is an essential macromineral for plants, which is studied extensively in edaphology in order to understand plant uptake from soil systems. In ecological terms, phosphorus is often a limiting factor in many environments; i.e. the availability of phosphorus governs the rate of growth of many organisms. In ecosystems an excess of phosphorus can be problematic, especially in aquatic systems, see eutrophication and algal blooms.

Precautions

Organic compounds of phosphorus form a wide class of materials, some of which are extremely toxic. Fluorophosphate esters are among the most potent neurotoxins known. A wide range of organophosphorus compounds are used for their toxicity to certain organisms as pesticides (herbicides, insecticides, fungicides, etc.) and weaponised as nerve agents. Most inorganic phosphates are relatively nontoxic and essential nutrients. For environmentally adverse effects of phosphates see eutrophication and algal blooms.

The white phosphorus allotrope should be kept under water at all times as it presents a significant fire hazard due to its extreme reactivity with atmospheric oxygen, and it should only be manipulated with forceps since contact with skin can cause severe burns. Chronic white phosphorus poisoning leads to necrosis of the jaw called "phossy jaw". Ingestion of white phosphorus may cause a medical condition known as "Smoking Stool Syndrome".

When the white form is exposed to sunlight or when it is heated in its own vapour to 250 °C, it is transmuted to the red form, which does not chemoluminesce in air. The red allotrope does not spontaneously ignite in air and is not as dangerous as the white form. Nevertheless, it should be handled with care because it reverts to white phosphorus in some temperature ranges and it also emits highly toxic fumes that consist of phosphorus oxides when it is heated.

Upon exposure to elemental phosphorus, in the past it was suggested to wash the affected area with 2% copper sulfate solution to form harmless compounds that can be washed away. According to the recent US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries, "Cupric (copper(II)) sulfate has been used by U.S. personnel in the past and is still being used by some nations. However, copper sulfate is toxic and its use will be discontinued. Copper sulfate may produce kidney and cerebral toxicity as well as intravascular hemolysis."

The manual suggests instead "a bicarbonate solution to neutralize phosphoric acid, which will then allow removal of visible white phosphorus. Particles often can be located by their emission of smoke when air strikes them, or by their phosphorescence in the dark. In dark surroundings, fragments are seen as luminescent spots." Then, "Promptly debride the burn if the patient's condition will permit removal of bits of WP which might be absorbed later and possibly produce systemic poisoning. DO NOT apply oily-based ointments until it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns." As white phosphorus readily mixes with oils, any oily substances or ointments are not recommended until the area is thoroughly cleaned and all white phosphorus removed.

US DEA List I status

Phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine. For this reason, two allotropes of elemental phosphorus—red phosphorus and white phosphorus—were designated by the United States Drug Enforcement Administration as List I precursor chemicals under 21 CFR 1310.02 effective on November 17, 2001. As a result, in the United States, handlers of red phosphorus or white phosphorus are subject to stringent regulatory controls pursuant to the Controlled Substances Act in order to reduce diversion of these substances for use in clandestine production of controlled substances.

See also

• White phosphorus

Notes

References

Notes

Sources

• Emsley, John (2000). The Shocking history of Phosphorus. A biography of the Devil's Element. London: MacMillan. ISBN 0-333-76638-5.
• Parkes, G.D. and Mellor, J.W. (1939). Mellor's Modern Inorganic Chemistry. London: Longman's Green and Co.
• Podger, Hugh (2002). Albright & Wilson. The Last 50 years. Studley: Brewin Books. ISBN 1-85858-223-7.
• Threlfall, Richard E. (1951). The Story of 100 years of Phosphorus Making: 1851–1951. Oldbury: Albright & Wilson ltd.

External links

• Further warnings of toxic effects and recommendations for treatment can be found in
• WebElements.com: Phosphorus
• eMedicine.com: Article on White Phophorus as used as weapon
• Website of the Technische Universität Darmstadt and the CEEP about Phosphorus Recovery