Phosphorus

This article is about the chemical element. For other uses, see Phosphorus (disambiguation).

silicon ← phosphorus → sulfur N ↑ P ↓ As 15P Periodic table
Appearance
colorless, waxy white, yellow, scarlet, red, violet, black waxy white (yellow cut), red (granules center left, chunk center right), and violet phosphorus
General properties
Name, symbol, number	phosphorus, P, 15
Pronunciation	/ˈfɒsfərəs/ fos-fər-əs
Element category	nonmetal
Group, period, block	15, 3, p
Standard atomic weight	30.973762(2)
Electron configuration	[Ne] 3s2 3p3
Electrons per shell	2, 8, 5 (Image)
Physical properties
Density (near r.t.)	(white) 1.823, (red) ≈ 2.2 – 2.34, (violet) 2.36, (black) 2.69 g·cm−3
Melting point	(white) 44.2 °C, (black) 610 °C
Sublimation point	(red) ≈ 416 – 590  °C, (violet) 620 °C
Boiling point	(white) 280.5 °C
Heat of fusion	(white) 0.66 kJ·mol−1
Heat of vaporization	(white) 12.4 kJ·mol−1
Molar heat capacity	(white) 23.824 J·mol−1·K−1
Vapor pressure (white)
P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 279 307 342 388 453 549
Vapor pressure (red, bp. 431 °C)
P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 455 489 529 576 635 704
Atomic properties
Oxidation states	5, 4, 3, 2[1], 1 [2], -1, -2, -3 (mildly acidic oxide)
Electronegativity	2.19 (Pauling scale)
Ionization energies (more)	1st: 1011.8 kJ·mol−1
	2nd: 1907 kJ·mol−1
	3rd: 2914.1 kJ·mol−1
Covalent radius	107±3 pm
Van der Waals radius	180 pm
Miscellanea
Magnetic ordering	(white,red,violet,black) diamagnetic[3]
Thermal conductivity	(white) 0.236, (black) 12.1 W·m−1·K−1
Bulk modulus	(white) 5, (red) 11 GPa
CAS registry number	7723-14-0
Most stable isotopes
Main article: Isotopes of phosphorus
iso NA half-life DM DE (MeV) DP 31P 100% 31P is stable with 16 neutrons 32P syn 14.28 d β− 1.709 32S 33P syn 25.3 d β− 0.249 33S
v · d · e · r

Phosphorus ( /ˈfɒsfərəs/ fos-fər-əs) is the chemical element that has the symbol P and atomic number 15. A multivalent nonmetal of the nitrogen group, phosphorus as a mineral is almost always present in its maximally oxidized state, as inorganic phosphate rocks. Elemental phosphorus exists in two major forms—white phosphorus and red phosphorus—but due to its high reactivity, phosphorus is never found as a free element on Earth.

The first form of elemental phosphorus to be produced (white phosphorus, in 1669) emits a faint glow upon exposure to oxygen – hence its name given from Greek mythology, Φωσφόρος meaning "light-bearer" (Latin Lucifer), referring to the "Morning Star", the planet Venus. Although the term "phosphorescence", meaning glow after illumination, derives from this property of phosphorus, the glow of phosphorus originates from oxidation of the white (but not red) phosphorus and should be called chemiluminescence.

The vast majority of phosphorus compounds are consumed as fertilizers. Other applications include the role of organophosphorus compounds in detergents, pesticides and nerve agents, and matches.^[4]

Phosphorus is essential for life. As phosphate, it is a component of DNA, RNA, ATP, and also the phospholipids that form all cell membranes. Demonstrating the link between phosphorus and life, elemental phosphorus was historically first isolated from human urine, and bone ash was an important early phosphate source. Phosphate minerals are fossils. Low phosphate levels are an important limit to growth in some aquatic systems. Today, the most important commercial use of phosphorus-based chemicals is the production of fertilizers, to replace the phosphorus that plants remove from the soil.

Characteristics

Main article: Allotropes of phosphorus

Physical

P₄ molecule

Phosphorus exists as several forms (allotropes) that exhibit strikingly different properties.^[5] The two most common allotropes are white phosphorus and red 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 standard atmospheres or 1.2 gigapascals). 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.^[6]

White phosphorus

White phosphorus exposed to air glows in the darkness

White phosphorus and related molecular forms

The most important form of elemental phosphorus from the perspective of applications and chemical literature is white phosphorus. It consists of tetrahedral P₄ molecules, in which each atom is bound to the other three atoms by a single bond. This P₄ tetrahedron is also present in liquid and gaseous phosphorus up to the temperature of 800 °C when it starts decomposing to P₂ molecules.^[7] Solid white exists in two forms. At low-temperatures, the β form is stable. At high-temperatures α form is predominant. These forms differ in terms of the relative orientations of the constituent P₄ tetrahedra.

White phosphorus is the least stable, the most reactive, more volatile, less dense, and more toxic than the other allotropes. White phosphorus gradually changes to red phosphorus. This transformation, which is accelerated by light and heat, and samples of white phosphorus almost always contain some red phosphorus and therefore appear yellow. For this reason, it is also called yellow phosphorus. It glows in the dark (when exposed to oxygen) with a very faint tint of green and blue, 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.^[8]

Thermolysis (cracking) of P₄ at 1100 kelvin) gives diphosphorus, P₂. This species not stable as a solid or liquid. The dimeric unit contains a triple bond and is analogous to N₂. It can also be generated as a transient intermediate in solution by thermolysis of organophosphorus precursor reagents.^[9] At still higher temperatures, P₂ dissociates into atomic P.

Although the term phosphorescence is derived from phosphorus, the reaction that 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).

Crystal structure of red phosphorus

red phosphorus

Red phosphorus

Red phosphorus is polymeric in structure. It can be viewed as a derivative of P₄ wherein one P-P bond is broken, and one additional bond is formed between the neighbouring tetrahedron resulting in a chain-like structure. Red phosphorus may be formed by heating white phosphorus to 250 °C (482 °F) or by exposing white phosphorus to sunlight.^[10] Phosphorus after this treatment is amorphous. Upon further heating, this material crystallises. In this sense, red phosphorus is not an 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 ignite in air at temperatures below 260 °C, whereas white phosphorus ignites at about 30 °C.^[11] It has commercial uses as a fire-retardant and in matches.

Violet phosphorus

Violet phosphorus is a form of phosphorus that can be produced by day-long annealing 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).^[6]

Crystal structure of black phosphorus

Black phosphorus

Black phosphorus is the least reactive allotrope and the thermodynamically stable form below 550 °C. It is also known as β-metallic phosphorus and has a structure somewhat resembling that of graphite.^[12][13] High pressures are usually required to produce black phosphorus, but it can also be produced at ambient conditions using metal salts as catalysts.^[14]

Properties of some allotropes of phosphorus^[5][6]

Form	white(α)	white(β)	violet	black
Symmetry	Body-centred cubic	Triclinic	Monoclinic	Orthorhombic
Pearson symbol		aP24	mP84	oS8
Space group	I43m	P1 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

Main article: Isotopes of phosphorus

Twenty-three isotopes of phosphorus are known,^[15] including all possibilities from ²⁴P up to ⁴⁶P). Only ³¹P is stable and is therefore present at 100% abundance. The half-integer nuclear spin and high abundance of ³¹P make phosphorus-31 NMR spectroscopy a very useful analytical tool in studies of phosphorus-containing samples.

Two radioactive isotopes of phosphorus have half-lives that 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 probes, 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 Bremsstrahlung (braking radiation). Therefore shielding must be accomplished with low density materials, e.g. Plexiglas (Lucite), other plastics, water, or (when transparency is not required), even wood.^[16]
• 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.

Occurrence

See also Category: Phosphate minerals.

In terms of stellar nucleosynthesis, stable forms of phosphorus are produced in large (greater than 3 solar masses) stars by fusing two oxygen atoms together.^{[citation needed]} This requires temperatures above 1,000 megakelvins.

Phosphorus is not found free in nature, but it is widely distributed in many minerals, mainly phosphates. 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.^[17] 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.^[18] 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.^[19]

Scarcity

Recent reports suggest that production of phosphorus may have peaked, leading to the possibility of global shortages by 2040.^[20] In 2007, at the rate of consumption, the supply of phosphorus was estimated to run out in 345 years.^[21] 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."^[22]

Production

Main article: phosphoric acid

The majority of phosphorus-containing compounds are produced for use as fertilisers. For this purpose, phosphate-containing minerals are converted to phosphoric acid. Two distinct routes are employed, the main one being treatment of phosphate minerals with sulfuric acid. The other process utilises white phosphorus, which may be produced by reaction and distillation from very low grade phosphate sources. The white phosphorus is then oxidised to phosphoric acid and subsequently neutralised with base to give phosphate salts. Phosphoric acid obtained via white phosphorus is relatively pure and is the main source of phosphates used in detergents and other non-fertiliser applications.

Elemental phosphorus

About 1,000,000 short tons (910,000 t) of elemental phosphorus is produced annually. 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 P₄. The product is subsequently condensed into a white powder under water to prevent oxidation by air. Even under water, white phosphorus is slowly converted to the more stable red phosphorus allotrope. The chemical equation for this process when starting with fluoroapatite, a common phosphate mineral, is:

4 Ca₅(PO₄)₃F + 18 SiO₂ + 30 C → 3 P₄ + 30 CO + 18 CaSiO₃ + 2 CaF₂

Side products from this production include ferrophosphorus, a crude form of Fe₂P, resulting from iron impurities in the mineral precursors. The silicate slag is a useful construction material. The fluoride is sometimes recovered for use in water fluoridation. More problematic is a "mud" containing significant amounts of white phosphorus. Production of white phosphorus is conducted in large facilities in part because it is energy intensive. The white phosphorus is transported in molten form. Some major accidents have occurred during transportation, train derailments at Brownston, Nebraska and Miamisburg, Ohio led to large fires. The worst incident in recent times was an environmental one in 1968 when the sea became contaminated due to spillages and/or inadequately treated sewage from a white phosphorus plant at Placentia Bay, Newfoundland.^[23]

Thermphos International is Europe's only producer of elemental phosphorus. About 6,500 short tons (5,900 t) of red phosphorus were produced in 1988. It is produced by heating white phosphorus near 270 °C. The product is treated to remove traces of white phosphorus. Red phosphorus is a component of fire-retardants and striking surfaces for matches.

Compounds

See also Category: Phosphorus compounds.

Oxyacids of phosphorus

Phosphorous oxyacids are extensive, often commercially important, and sometimes structurally complicated. They all have acidic protons bound to oxygen atoms and some have nonacidic protons that are bonded directly to phosphorus. Although many oxyacids of phosphorus are formed, only six are important, 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

Phosphorus(V) compounds

Oxides

The most prevalent compounds of phosphorus are derivatives of phosphate (PO₄^3-), a tetrahedral anion.^[24] Phosphate is the conjugate base of phosphoric acid, which is produced on a massive scale for use in fertilizers. Being triprotic, phosphoric acid converts stepwise to three conjugate bases:

H₃PO₄ + H₂O H₃O⁺ + H₂PO₄⁻       K_a1= 7.25×10⁻³

H₂PO₄⁻ + H₂O H₃O⁺ + HPO₄²⁻       K_a2= 6.31×10⁻⁸

HPO₄²⁻ + H₂O H₃O⁺ +  PO₄³⁻        K_a3= 3.98×10⁻¹³

Phosphate exhibits the tendency to form chains and rings with P-O-P bonds. Many polyphosphates are known, including ATP. Polyphosphates arise by dehydration of hydrogen phosphates such as HPO₄^2- and H₂PO₄^-. For example, the industrially important trisodium triphosphate (also known as sodium tripolyphosphate, STPP) is produced industrially on a megatonne scale via this condensation reaction:

2 Na₂[(HO)PO₃] + Na[(HO)₂PO₂] → Na₅[O₃P-O-P(O)₂-O-PO₃] + 2 H₂O

Phosphorus pentoxide (P₄O₁₀) is the acid anhydride of phosphoric acid, but several intermediates are known between the two. This waxy white solid reacts vigorously with water.

The tetrahedral structure of P₄O₁₀ and P₄S₁₀.

With metal cations, phosphate forms a variety of salts. These solids are polymeric, featuring P-O-M linkages. When the metal cation has a charge of 2+ or 3+, the salts are generally insoluble, hence they exist as common minerals. Many phosphate salts are derived from hydrogen phosphate (HPO₄^2-).

PCl₅ and PF₅ are common compounds. Both are volatile and pale or colourless. The other two halides, PBr₅ and PI₅PI₅ are unstable. The pentachloride and pentafluoride adopt trigonal bipyramid molecular geometry and are Lewis acids. With fluoride, PF₅ forms PF₆^–, an anion that is isoelectronic with SF₆. The most important oxyhalide is phosphorus oxychloride (POCl₃), which is tetrahedral.

Before extensive computer calculations were feasible, it was proposed that bonding in phosphorus(V) compounds involved d orbitals. It is now accepted that the bonding can be better explained by molecular orbital theory and involves only s- and p-orbitals on phosphorus.^[25]

Nitrides

Compounds of the formula (PNCl₂)_n exist mainly as rings such as the trimer hexachlorophosphazene. The phosphazenes arise by treatment of phosphorus pentachloride with ammonium chloride: PCl₅ + NH₄Cl → 1/n (NPCl₂)_n + 4 HCl The chloride groups can be replaced by alkoxide (RO^-) to give rise to a family of polymers with potentially useful properties.^[26]

Sulfides

Main article: phosphorus sulfide

Phosphorus forms a wide range of sulfides, where phosphorus can be P(V), P(III) or other oxidation states. Most famous is the three-fold symmetric P4S3 used in strike-anywhere matches. P₄S₁₀ and P₄O₁₀ have analogous structures.^[27]

Phosphorus(III) compounds

Phosphine (PH₃) and its organic derivatives (PR₃) are structural analogues with ammonia (NH₃) but the angles at phosphorus are closer to 90° for phosphine and its organic derivatives. It is an ill-smelling, toxic compound. Phosphine is produced by hydrolysis of calcium phosphide, Ca₃P₂. Unlike ammonia, phosphine is oxidised by air. Phosphine is also far less basic than ammonia.

All four symmetrical trihalides are well known: gaseous PF₃, the yellowish liquids PCl₃ and PBr₃, and the solid PI₃. These materials are moisture sensitive, hydrolysing to give phosphorus acid. The trichloride, a common reagent, is produced by chlorination of white phosphorus:

P₄ + 6 Cl₂ → 4 PCl₃

The trifluoride is produced by from the trichloride by halide exchange. PF₃ is toxic because it binds to haemoglobin.

Phosphorus(III) oxide, P₄O₆ (also called tetraphosphorus hexoxide) is the anhydride of P(OH)₃, the minor tautomer of phosphorous acid. The structure of P₄O₆ is like that of P₄O₁₀ less the terminal oxide groups.

Mixed oxyhalides and oxyhydrides of phosphorus(III) are almost unknown.

Organophosphorus compounds

Main article: organophosphorus compounds

Compounds with P-C and P-O-C bonds are often classified as organophosphorus compounds. They are widely used commercially. The PCl₃ serves as a source of P³⁺ in routes to organophosphorus(III) compounds. For example it is the precursor to triphenylphosphine:

PCl₃ + 6 Na + 3 C₆H₅Cl → P(C₆H₅)₃ + 6 NaCl

Treatment of phosphorus trihalides with alcohols and phenols gives phosphites, e.g. triphenylphosphite:

PCl₃ + 3 C₆H₅OH → P(OC₆H₅)₃ + 3 HCl

Similar reactions occur for phosphorus oxychloride, affording triphenylphosphate:

OPCl₃ + 3 C₆H₅OH → OP(OC₆H₅)₃ + 3 HCl

Phosphorus(I) and phosphorus(II) compounds

These compounds generally feature P-P bonds.^[8] Examples include catenated derivatives of phosphine and organophosphines. The highly flammable gas diphosphine (P₂H₄) is the first of a series of derivatives of this type. Diphosphine is an analogue of hydrazine. Compounds containing P=P double bonds have also been observed, although they are rare.

A stable diphosphene, a derivative of phosphorus(I).

Phosphides

The phosphide ion is P^3-. Phosphides arise by reaction of metals with red phosphorus. Salts of P^3- do not exist in solution and these derivatives are refractory, reflecting their high lattice energy. Illustrated by the behaviour calcium phosphide, many metal phosphides hydrolyse in water with release of phosphine:

Ca₃P₂ + 6 H₂O → 2 PH₃ + 3 Ca(OH)₂

Schreibersite is a naturally occurring phosphide found in meteorites. Many polyphosphides are known such as derivatives of OsP₂. These can be structurally complex ranging from Na₃P₇ and derivatives of P₂₆^4-. Often these species adopt cage-like structures that resemble fragments of violet phosphorus.

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, φέρω = carry), which roughly translates as light-bringer or light carrier.^[10] (In Greek mythology and tradition, Augerinus (Αυγερινός = morning star, in use until today), Hesperus or Hesperinus (΄Εσπερος or Εσπερινός or Αποσπερίτης = evening star, in use until today) and Eosphorus (Εωσφόρος = dawnbearer, not in use for the planet after Christianity) 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 (e.g., phosphorous acid) and P⁵⁺ valence phosphoric compounds (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.^[28] Brand experimented with urine, which contains considerable quantities of dissolved phosphates from normal metabolism.^[10] 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. It was named phosphorus mirabilis ("miraculous bearer of light").^[29] 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.

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;^[30] there is only a range of partial pressure at which it does. Heat can be applied to drive the reaction at higher pressures.^[31]

In 1974, the glow was explained by R. J. van Zee and A. U. Khan.^[32] 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.

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.

Brand at first tried to keep the method secret,^[33] but later sold the recipe for 200 thalers to D Krafft from Dresden,^[10] 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.^[10] 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.^[34] 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.

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.^[35] 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 when the submerged-arc furnace for phosphorus production was introduced to reduce phosphate rock.^[36][37] The electric furnace method allowed production to increase to the point where phosphorus could be used in weapons of war.^[32][38] In World War I it was used in incendiaries, smoke screens and tracer bullets.^[38] A special incendiary bullet was developed to shoot at hydrogen-filled Zeppelins over Britain (hydrogen being highly flammable if it can be ignited).^[38] 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.^[8]

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).^[32] 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 (1906), requiring its adoption as a safer alternative for match manufacture.^[39] The toxicity of white phosphorus led to discontinuation of it use in matches.^[40] Ironically, the Allies used phosphorus incendiary bombs in World War II to destroy Hamburg, the place where the "miraculous bearer of light" was first discovered.^[29]

Applications

Fertiliser

The dominant application of phosphorus is in fertilisers, which provides phosphate as required for all life and is often a limiting nutrient for crops. 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 fertilisers 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 fertilisers that 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.

Match striking surface made of a mixture of red phosphorus, glue and ground glass. The glass powder is used to increase the friction.

Widely used compounds	Use
Ca(H2PO4)2·H2O	Baking powder and 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

Organophosphorus compounds

White phosphorus is widely used to make organophosphorus compounds, through the intermediates phosphorus chlorides and two phosphorus sulfides, phosphorus pentasulfide, and phosphorus sesquisulfide.^[41] Organophosphorus compounds have many applications, including in plasticizers, flame retardants, pesticides, extraction agents, and water treatment.^[8] in particular the herbicide glyphosate sold under the brand name Roundup.

Metallurgical aspects

Phosphorus is also an important component in steel production, in the making of phosphor bronze, and in many other related products. Phosphorus is added to metallic copper during its smelting process to react with oxygen present as an impurity in copper and to produce oxygen-free copper or phosphorus-containing copper (CuOFP) alloys with a higher thermal and electrical conductivity than normal copper.

Matches

Phosphorus sesquisulfide (P₄S₃) is used in heads of strike-anywhere matches. Red phosphorus is a component of striking surface for matches and flares. It is used in cap gun caps.^[41]

Water softening

Sodium tripolyphosphate made from phosphoric acid is used in laundry detergents in some countries, but banned for this use in others.^[42] It is useful for softening water to enhance the performance of the detergents and to prevent pipe/boiler tube corrosion.

Niche applications

• Phosphates are utilized in the making of special glasses that are used for sodium lamps.^[42]
• Bone-ash, calcium phosphate, is used in the production of fine china.^[42]
• 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.^[41] These include mono-calcium phosphate that is employed in baking powder and sodium tripolyphosphate and other sodium phosphates.^[41] Among other uses these are used to improve the characteristics of processed meat and cheese. Others are used in toothpaste.^[41]
• White phosphorus, called "WP" (slang term "Willie Peter") is used in military applications as incendiary bombs, 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.^[43][44]
• In trace amounts, phosphorus is used as a dopant for n-type semiconductors.
• ³²P and ³³P are used as radioactive tracers in biochemical laboratories.
• Phosphate is a strong complexing agent for the hexavalent uranyl (UO₂²⁺) species and this is the reason why apatite and other natural phosphates can often be very rich in uranium.
• Tributylphosphate is an organophosphate soluble in kerosene and used to extract uranium in the Purex process applied in the reprocessing of spent nuclear fuel.

Biological role

Inorganic phosphorus in the form of the phosphate PO₄^3– is required for all known forms of life,^[45] playing 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.^[8]

Living cells are defined by a membrane that separates it from its surroundings. 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.^[46]

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.^[47] 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 phosphate ions such as H₂PO₄^- and HPO₄^2-. Only about 0.1% of body phosphate circulates in the blood, and this amount reflects the amount of phosphate available to soft tissue cells.

Bone and teeth enamel

The main component of bone is hydroxyapatite as well as amorphous forms of calcium phosphate, possibly including carbonate. Hydroxyapatite is the main component of tooth enamel. Water fluoridation enhances the resistance of teeth to decay by the partial conversion of this mineral to the still harder material called fluoroapatite:^[8]

Ca₅(PO₄)₃OH + F^- → Ca₅(PO₄)₃F + OH^-

Phosphorus deficiency

In medicine, low-phosphate syndromes are caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes that 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, 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.^[48]

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, resulting in eutrophication which sometimes lead to algal blooms.

Food sources

The main food sources of phosphorus are kinds of food containing protein, for example, milk and meat. As a rule, if one holds a meal plan providing sufficient amount of protein and calcium then amount of phosphorus is also sufficient.^[49]

Precautions

Organic compounds of phosphorus form a wide class of materials, many are required for life, but some 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.^[8]

The white phosphorus allotrope presents a significant hazard because it ignites in air and produces phosphoric acid residue. 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".^[50]

Phosphorus explosion

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."^[51]

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. Promptly debride the burn if the patient's condition will permit removal of bits of WP (white phosphorus) that 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."^{[note 1]} 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.^[52] For this reason, red 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.^[53] As a result, in the United States, handlers of red or white phosphorus are subject to stringent regulatory controls.^[53][54][55]

See also

• White phosphorus (weapon)
• Peak phosphorus

Notes

1. ^ This quote uses the word "phosphorescence", which is incorrect; WP, (white phosphorus), exhibits chemoluminescence upon exposure to air and if there is any WP in the wound, covered by tissue or fluids such as blood serum, it will not chemoluminesce until it is moved to a position where the air can get at it and activate the chemoluminescent glow, which requires a very dark room and dark adapted eyes to see clearly.

References

Notes

1. ^ webelements
2. ^ Ellis, Bobby D.; MacDonald, Charles L. B. (2006). "Phosphorus(I) Iodide:  A Versatile Metathesis Reagent for the Synthesis of Low Oxidation State Phosphorus Compounds". Inorganic Chemistry 45 (17): 6864–74. doi:10.1021/ic060186o. PMID 16903744. 
3. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81st edition, CRC press.
4. ^ Herbert Diskowski, Thomas Hofmann "Phosphorus" in Ullmann's Encyclopedia of Industrial Chemistry 2005, Wiley-VCH, Weinheim. doi:10.1002/14356007.a19_505
5. ^ ^{a b} A. Holleman, N. Wiberg (1985). "XV 2.1.3". Lehrbuch der Anorganischen Chemie (33 ed.). de Gruyter. ISBN 3110126419. 
6. ^ ^{a b c} Berger, L. I. (1996). Semiconductor materials. CRC Press. p. 84. ISBN 0849389127. http://books.google.com/?id=Ty5Ymlg_Mh0C&pg=PA84. 
7. ^ Simon, Arndt; Borrmann, Horst; Horakh, Jörg (1997). "On the Polymorphism of White Phosphorus". Chemische Berichte 130: 1235. doi:10.1002/cber.19971300911. 
8. ^ ^{a b c d e f g} Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
9. ^ Piro, N. A.; Figueroa, JS; McKellar, JT; Cummins, CC (2006). "Triple-Bond Reactivity of Diphosphorus Molecules". Science 313 (5791): 1276. Bibcode 2006Sci...313.1276P. doi:10.1126/science.1129630. PMID 16946068. 
10. ^ ^{a b c d e} Parkes & Mellor 1939, p. 717
11. ^ Parkes & Mellor 1939, pp. 721–722
12. ^ A. Brown, S. Runquist (1965). "Refinement of the crystal structure of black phosphorus". Acta Crystallogr. 19: 684. doi:10.1107/S0365110X65004140. 
13. ^ Cartz, L.;Srinivasa, S.R.;Riedner, R.J.;Jorgensen, J.D.;Worlton, T.G. (1979). "Effect of pressure on bonding in black phosphorus". Journal of Chemical Physics 71: 1718–1721. Bibcode 1979JChPh..71.1718C. doi:10.1063/1.438523. 
14. ^ Lange, Stefan; Schmidt, Peer and Nilges, Tom (2007). "Au3SnP7@Black Phosphorus: An Easy Access to Black Phosphorus". Inorg. Chem. 46 (10): 4028. doi:10.1021/ic062192q. PMID 17439206. 
15. ^ "The Berkeley Laboratory Isotopes Project". http://ie.lbl.gov/education/parent/P_iso.htm. Retrieved 2009-05-05. 
16. ^ "Phosphorus-32". University of Michigan Department of Occupational Safety & Environmental Health. http://www.oseh.umich.edu/pdf/TrainP32.pdf. Retrieved 2010-11-18. 
17. ^ "Phosphate Rock: Statistics and Information". USGS. http://minerals.usgs.gov/minerals/pubs/commodity/phosphate_rock/. Retrieved 2009-06-06. 
18. ^ Threlfall 1951, p. 51
19. ^ Podger 2002, pp. 297–298
20. ^ Carpenter S.R. and Bennett E.M. (2011). "Reconsideration of the planetary boundary for phosphorus". Environmental Research Letters 6 (1): 1–12. Bibcode 2011ERL.....6a4009C. doi:10.1088/1748-9326/6/1/014009. 
21. ^ "How Long Will it Last?". New Scientist 194 (2605): 38–39. May 26, 2007. Bibcode 2007NewSc.194...38R. doi:10.1016/S0262-4079(07)61508-5. ISSN 4079 0262 4079. 
22. ^ Lewis, Leo (2008-06-23). "Scientists warn of lack of vital phosphorus as biofuels raise demand". The Times. http://business.timesonline.co.uk/tol/business/industry_sectors/natural_resources/article4193017.ece. 
23. ^ "ERCO and Long Harbour". Memorial University of Newfoundland and the C.R.B. Foundation. http://www.heritage.nf.ca/law/erco.html. Retrieved 2009-06-06. 
24. ^ D. E. C. Corbridge "Phosphorus: An Outline of its Chemistry, Biochemistry, and Technology" 5th Edition Elsevier: Amsterdam 1995. ISBN 0-444-89307-5.
25. ^ Kutzelnigg, W. (1984). "Chemical Bonding in Higher Main Group Elements". Angewandte Chemie Int. (English) Ed. 23: 272–295. doi:10.1002/anie.198402721. http://web.uvic.ca/~chem421/ACIE_1984_Kutzelnigg_review.pdf. 
26. ^ Mark, J. E.; Allcock, H. R.; West, R. “Inorganic Polymers” Prentice Hall, Englewood, NJ: 1992. ISBN 0-13-465881-7.
27. ^ Heal, H. G. “The Inorganic Heterocyclic Chemistry of Sulfur, Nitrogen, and Phosphorus” Academic Press: London; 1980. ISBN 0-12-33580-6.
28. ^ Beatty, Richard (2000). Phosphorus. Marshall Cavendish. p. 7. ISBN 0761409467. http://books.google.com/?id=FHJIUJM1_JUC&pg=PA7. 
29. ^ ^{a b} "Experts Warn of Impending Phosphorus Crisis", by Hilmar Schmundt, Spiegel, 21 April 2010
30. ^ "Nobel Prize in Chemistry 1956 – Presentation Speech by Professor A. Ölander (committee member)". http://nobelprize.org/chemistry/laureates/1956/press.html. Retrieved 2009-05-05. 
31. ^ "Phosphorus Topics page, at Lateral Science". http://www.lateralscience.co.uk/phos/index.html. Retrieved 2009-05-05. 
32. ^ ^{a b c} Emsley, John (2000). The Shocking History of Phosphorus. London: Macmillan. ISBN 0-330-39005-8. 
33. ^ J. M. Stillman (1960). The Story of Alchemy and Early Chemistry. New York: Dover. pp. 418–419. ISBN 0766132307. 
34. ^ cf. "Memoir on Combustion in General" Mémoires de l'Académie Royale des Sciences 1777, 592–600. from Henry Marshall Leicester and Herbert S. Klickstein, A Source Book in Chemistry 1400–1900 (New York: McGraw Hill, 1952)
35. ^ Threlfall 1951, pp. 49–66
36. ^ Threlfall 1951, pp. 81–101
37. ^ Parkes & Mellor 1939, p. 718–720.
38. ^ ^{a b c} Threlfall 1951, pp. 167–185
39. ^ Lewis R. Goldfrank, Neal Flomenbaum, Mary Ann Howland, Robert S. Hoffman, Neal A. Lewin, Lewis S. Nelson (2006). Goldfrank's toxicologic emergencies. McGraw-Hill Professional. pp. 1486–1489. ISBN 0071437630. http://books.google.com/?id=cvJuLqBxGUcC&pg=PA1487. 
40. ^ The White Phosphorus Matches Prohibition Act, 1908.
41. ^ ^{a b c d e} Threlfall, R.E. (1951). 100 years of Phosphorus Making: 1851–1951. Oldbury: Albright and Wilson Ltd. 
42. ^ ^{a b c} Hammond, C. R. (2000). The Elements, in Handbook of Chemistry and Physics 81st edition. CRC press. ISBN 0849304814. 
43. ^ "Obsolete hand grenades". GlobalSecurity.Org. http://www.globalsecurity.org/military/library/policy/army/fm/3-23-30/appe.htm. Retrieved 2009-08-03. 
44. ^ Dockery, Kevin (1997). Special Warfare Special Weapons. Chicago: Emperor's Press. ISBN 1-883-47600-3. 
45. ^ Phosphorus Cycle – Terrestrial Phosphorus Cycle, Transport of Phosphorus, from Continents to the Ocean, The Marine Phosphorus Cycle
46. ^ Nelson, D. L.; Cox, M. M. "Lehninger, Principles of Biochemistry" 3rd Ed. Worth Publishing: New York, 2000. ISBN 1-57259-153-6.
47. ^ Bernhardt, Nancy E.; Kasko, Artur M. (2008). Nutrition for the Middle Aged and Elderly. Nova Publishers. p. 171. ISBN 1604561467. http://books.google.com/?id=ba_5OSsyS4YC&pg=PA171. 
48. ^ Anderson, John J. B. (Apr 1996). "Calcium, Phosphorus and Human Bone Development". Journal of Nutrition 126 (4 Suppl.): 1153S–1158S. PMID 8642449. http://jn.nutrition.org/cgi/content/abstract/126/4_Suppl/1153S. 
49. ^ http://www.nlm.nih.gov/medlineplus/ency/article/002424.htm
50. ^ "CBRNE – Incendiary Agents, White Phosphorus (Smoking Stool Syndrome)". http://www.emedicine.com/EMERG/topic918.htm. Retrieved 2009-05-05. 
51. ^ "US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries". http://www.vnh.org/FM8285/Chapter/chapter9.html. Retrieved 2009-05-05. ^{[dead link]}
52. ^ Skinner, H.F. (1990). "Methamphetamine synthesis via hydriodic acid/red phosphorus reduction of ephedrine". Forensic Science International 48 (2): 123–134. doi:10.1016/0379-0738(90)90104-7. 
53. ^ ^{a b} "66 FR 52670—52675". 17 October 2001. http://frwebgate.access.gpo.gov/cgi-bin/getdoc.cgi?dbname=2001_register&docid=01-26013-filed. Retrieved 2009-05-05. 
54. ^ "21 CFR 1309". http://www.access.gpo.gov/nara/cfr/waisidx_06/21cfr1309_06.html. Retrieved 2009-05-05. 
55. ^ "21 USC, Chapter 13 (Controlled Substances Act)". http://www.usdoj.gov/dea/pubs/csa.html. Retrieved 2009-05-05. 

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.; Mellor, J.W. (1939). Mellor's Modern Inorganic Chemistry. 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

• The Periodic Table of Videos video of Phosphorus at YouTube
• Further warnings of toxic effects and recommendations for treatment can be found in "Emergency War Surgery NATO Handbook: Part I: Types of Wounds and Injuries: Chapter III: Burn Injury: Chemical Burns And White Phosphorus injury". http://www.brooksidepress.org/Products/OperationalMedicine/DATA/operationalmed/Manuals/NATOEWS/ch03/03ChemicalBurns.html. Retrieved 2009-05-05. 
• WebElements.com: Phosphorus
• Simon, Fa; Pickering, Lk (Mar 1976). "Acute yellow phosphorus poisoning. "Smoking stool syndrome".". JAMA : the journal of the American Medical Association 235 (13): 1343–4. doi:10.1001/jama.235.13.1343. ISSN 0098-7484. PMID 946251. http://toxnet.nlm.nih.gov/cgi-bin/sis/search/r?dbs+hsdb:@term+@rn+7723-14-0. 
• eMedicine.com: Article on White Phophorus as used as weapon
• Website of the Technische Universität Darmstadt and the CEEP about Phosphorus Recovery