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siliconphosphorussulfur
N

P

As
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15P
Appearance
colorless, waxy white, yellow, scarlet, red, black
File:PhosphComby.jpg
General
Name, symbol, number phosphorus, P, 15
Element categorynonmetal
Group, period, block 153, p
Standard atomic weight 30.973762(2)g·mol−1
Electron configuration [Ne] 3s2 3p3
Electrons per shell 2, 8, 5 (Image)
Physical properties
Density (near r.t.)(white) 1.823, (red) 2.34, (black) 2.69 g·cm−3
Melting point(white) 44 °C, (red) 579 °C, (black) 610 °C
boiling point(white) 280 °C, (red: sublimes) 416 °C
Heat of fusion(white) 0.66 kJ·mol−1
Heat of vaporization12.4 kJ·mol−1
Specific heat capacity(25 °C) (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)
P/Pa 1 10 100 1 k 10 k 100 k
at T/K 455 489 529 576 635 704
Atomic properties

Oxidation states5, 4, 3, 2[1], 1 [2], -1, -2, -3
(mildly acidic oxide)
Electronegativity2.19 (Pauling scale)
Ionization energies
(more)
1st: 1011.8 kJ·mol−1
2nd: 1907 kJ·mol−1
3rd: 2914.1 kJ·mol−1
Covalent radius107±3 pm
Van der Waals radius180 pm
Miscellaneous
Magnetic orderingdiamagnetic[3]
Thermal conductivity(300 K) (white) 0.236, (black) 12.1 W·m−1·K−1
Bulk modulus11 GPa
CAS registry number7723-14-0
Most stable isotopes
Main article: Isotopes of phosphorus
iso N.A. 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
References

Phosphorus (pronounced /ˈfɒsfərəs/) 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, discovered 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.

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.

Contents

Physical properties

Glow from white phosphorus

Phosphorus was discovered by German alchemist Hennig Brand in 1674 or 1675. Working in Hamburg, Brand attempted to distill some kind of "life essence" from his urine, and in the process produced a white material that glowed in the dark.[4] 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;[5] there is only a range of partial pressure at which it does. Heat can be applied to drive the reaction at higher pressures.[6]

In 1974, the glow was explained by R. J. van Zee and A. U. Khan.[7] 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 (color) of the light emitted is proportional to the energy difference of the two electronic states involved.[8]

Allotropes

Phosphorus has several forms (allotropes) which have strikingly different properties. 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 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.[9]

[[File:|thumb|100px|P4O10 molecule]] White phosphorus has two forms, low-temperature β form and high-temperature α form. They both contain a phosphorus P4 tetrahedron as a structural unit, 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.[10] 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 odor 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.[11]

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.[12] 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.

[[File:|thumb|Crystal structure of red phosphorus]] In the red phosphorus, one of the P4 bonds is broken, and one additional bond is formed with a neighboring 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.[4] 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. Red phosphorus does not catch fire in air at temperatures below 260 °C, whereas white phosphorus ignites at about 30 °C.[13] In 1865, Hittorf discovered that when phosphorus was recrystallized from molten lead, a red/purple form is obtained. This purple form is sometimes known as "Hittorf's phosphorus" (or violet or α-metallic phosphorus).[9]

[[File:|thumb|upright|Crystal structure of black phosphorus]] Black phosphorus is the least reactive allotrope. It is also known as β-metallic phosphorus and has a structure somewhat resembling that of graphite.[14][15] High pressures are usually required to produce black phosphorus, but it can also be produced at ambient conditions using metal salts as catalysts.[16]

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).[17]

Properties of some allotropes of phosphorus[9]
Form white(α) white(β) red black
Symmetry Body-centered 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 2.34 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[18] (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 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 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.[19]
  • 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

See also Phosphorus compounds

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 hybridization, but nitrogen did not). However, in the early eighties the German theoretical chemist Werner Kutzelnigg[20] 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-center-4-electron (3c-4e) 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 analogs with ammonia (NH3) and form pyramidal structures with the phosphorus or arsenic atom in the center bound to three hydrogen atoms and one lone electron pair. Both are colorless, 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), analogs of ammonium salts, but these salts immediately decompose in water and do not yield phosphonium (PH4+) ions. Diphosphine (P2H4 or H2P-PH2) is an analog of hydrazine (N2H4) that is a colorless 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 halide. 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 halide 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(III) oxide, P4O6 (also called tetraphosphorus hexoxide) and phosphorus(IV) 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 StateFormulaNameAcidic ProtonsCompounds
+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

History and discovery

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.[4] (In Greek mythology, Hesperus (evening star) and Eosphorus (dawnbearer) are close homologues, and also associated with Phosphorus-the-planet).

The first recorded production of elemental phosphorus was in 1674 or 1675 by the German alchemist Hennig Brand through a preparation of urine, which contains considerable quantities of dissolved phosphates from normal metabolism.[4] 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,[21] but later sold the recipe for 200 thaler to D Krafft from Dresden,[4] 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.[4] 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).[7] In addition, exposure to the vapours gave match workers a 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. [11]

Spelling and etymology

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).

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.[22] 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.[12] 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.[23]

In 2007, at the current rate of consumption, the supply of phosphorus was estimated to run out in 345 years.[24] 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."[25]

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.[26] 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.[27][28] 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.[7][12] In World War I it was used in incendiaries, smoke screens and tracer bullets.[12] A special incendiary bullet was developed to shoot at hydrogen-filled Zeppelins over Britain (hydrogen being highly inflammable if it can be ignited).[12] 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).[11]

Today phosphorus production is larger than ever. It is used as a precursor for various chemicals,[29] 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.[30]

Applications

Widely used compoundsUse
Ca(H2PO4)2•H2OBaking powder & fertilizers
CaHPO4•2H2OAnimal food additive, toothpowder
H3PO4Manufacture of phosphate fertilizers
PCl3Manufacture of POCl3 and pesticides
POCl3Manufacturing plasticizer
P4S10Manufacturing of additives and pesticides
Na5P3O10Detergents

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 (PO43-) 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.

Biological role

Phosphorus is a key element in all known forms of life. Inorganic phosphorus in the form of the phosphate PO43- 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 via 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. [11]

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. [11]

An average adult human contains a little less than 1 kg of phosphorus, about 85% of which is present in bones and teeth in the form of apatite, and the remainder inside cells in soft tissues. 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. [11]

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 diarrhea and calcification (hardening) of organs and soft tissue, and can interfere with the body's ability to use iron, calcium, magnesium, and zinc.[33]

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 weaponized as nerve agents. Most inorganic phosphates are relatively nontoxic and essential nutrients. For environmentally adverse effects of phosphates see eutrophication and algal blooms. [11]

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". [34]

When the white form is exposed to sunlight or when it is heated in its own vapor 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. [11]


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."[35]

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."[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.[36] 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.[37] 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.[37][38][39]

See also

Notes

  1. ^ This quote uses the word "phosphorescence", which is actually incorrect, WP, (White Phosphorous), 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 chemoluminescence 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. ^ B. D. Ellis and C. L. B. Macdonald* "Phosphorus(I) Iodide: A Versatile Metathesis Reagent for the Synthesis of Low Oxidation State Phosphorus Compounds" Inorg. Chem., 2006, 45 (17), pp 6864
  3. ^ Magnetic susceptibility of the elements and inorganic compounds, in Handbook of Chemistry and Physics 81th edition, CRC press.
  4. ^ a b c d e f Mellor (1939), page 717.
  5. ^ "Nobel Prize in Chemistry 1956 - Presentation Speech by Professor A. Ölander (committee member)". http://nobelprize.org/chemistry/laureates/1956/press.html. Retrieved on 2009-05-05. 
  6. ^ "Phosphorus Topics page, at Lateral Science". http://www.lateralscience.co.uk/phos/index.html. Retrieved on 2009-05-05. 
  7. ^ a b c Emsley, John (2000). The Shocking History of Phosphorus. London: Macmillan. ISBN 0-330-39005-8. 
  8. ^ Ana M. García-Campaña, Willy R. G. Baeyens (2001). Chemiluminescence in analytical chemistry. CRC Press. pp. 2–12. ISBN 0824704649. http://books.google.com/books?id=-EPSISOfCxYC. 
  9. ^ a b c L. I. Berger (1996). Semiconductor materials. CRC Press. p. 84. ISBN 0849389127. http://books.google.com/books?id=Ty5Ymlg_Mh0C&pg=PA84. 
  10. ^ Simon, Arndt (1997). "On the Polymorphism of White Phosphorus". Chemische Berichte 130: 1235. doi:10.1002/cber.19971300911. 
  11. ^ a b c d e f g h i 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/books?id=cvJuLqBxGUcC&pg=PA1487. 
  12. ^ a b c d e f g h i j k Threlfall, R.E. (1951). 100 years of Phosphorus Making: 1851 - 1951. Oldbury: Albright and Wilson Ltd. 
  13. ^ Mellor (1939), pages 721–722.
  14. ^ A. Brown, S. Runquist (1965). "Refinement of the crystal structure of black phosphorus". Acta Crystallogr. 19: 684. doi:10.1107/S0365110X65004140. 
  15. ^ 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. doi:10.1063/1.438523. 
  16. ^ Stefan Lange, Peer Schmidt, and Tom Nilges (2007). "Au3SnP7@Black Phosphorus: An Easy Access to Black Phosphorus". Inorg. Chem. 46: 4028. doi:10.1021/ic062192q. 
  17. ^ Piro, N. A. (2006). "Triple-Bond Reactivity of Diphosphorus Molecules". Science 313 (5791): 1276. doi:10.1126/science.1129630. PMID 16946068. 
  18. ^ "The Berkeley Laboratory Isotopes Project". http://ie.lbl.gov/education/parent/P_iso.htm. Retrieved on 2009-05-05. 
  19. ^ "Occupational Safety & Environmental Health: Phsophorus-21". http://www.oseh.umich.edu/TrainP32.pdf. Retrieved on 2009-05-05. 
  20. ^ W. Kutzelnigg (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. 
  21. ^ J. M. Stillman (1960). The Story of Alchemy and Early Chemistry. New York: Dover. pp. 418–419. 
  22. ^ "Phosphate Rock: Statistics and Information". USGS. http://minerals.usgs.gov/minerals/pubs/commodity/phosphate_rock/. Retrieved on 2009-06-06. 
  23. ^ Podger (2002), pages 297–298.
  24. ^ "How Long Will it Last?". New Scientist 194 (2605): 38–39. May 26, 2007. ISSN 4079 0262 4079. 
  25. ^ Leo Lewis (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. 
  26. ^ Threlfall (1951), Chapter V: The foundations:1844–56: The phosphorus retort.
  27. ^ Threlfall (1951), Chapter VII: The second generation:1880–1915: Part 1: The electric furnace.
  28. ^ Mellor (1939), pages 718–720.
  29. ^ Aall C. H. (1952). "The American Phosphorus Industry". Industrial & Engineering Chemistry 44 (7): 1520–1525. doi:10.1021/ie50511a018. 
  30. ^ "ERCO and Long Harbour". Memorial University of Newfoundland and the C.R.B. Foundation. http://www.heritage.nf.ca/law/erco.html. Retrieved on 2009-06-06. 
  31. ^ a b c C. R. Hammond (2000). The Elements, in Handbook of Chemistry and Physics 81th edition. CRC press. ISBN 0849304814. 
  32. ^ Dockery, Kevin. (1997). Special Warfare Special Weapons. Chicago: Emperor's Press. ISBN 1-883-47600-3. 
  33. ^ Anderson, John J. B. (01 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. 
  34. ^ "CBRNE - Incendiary Agents, White Phosphorus (Smoking Stool Syndrome)". http://www.emedicine.com/EMERG/topic918.htm. Retrieved on 2009-05-05. 
  35. ^ "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 on 2009-05-05. 
  36. ^ 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. 
  37. ^ 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 on 2009-05-05. 
  38. ^ "21 CFR 1309". http://www.access.gpo.gov/nara/cfr/waisidx_06/21cfr1309_06.html. Retrieved on 2009-05-05. 
  39. ^ "21 USC, Chapter 13 (Controlled Substances Act)". http://www.usdoj.gov/dea/pubs/csa.html. Retrieved on 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. 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


Wiktionary

Up to date as of January 15, 2010

Definition from Wiktionary, a free dictionary

See also Phosphorus

Contents

English

Chemical element
P Previous: silicon (Si)
Next: sulfur (S)

Pronunciation

Etymology

From Latin phosphorus, from Ancient Greek φωσφόρος (fosfóros), the bearer of light), from φῶς (fos), light) + φέρω (phéro), to bear, to carry).

Noun

Singular
phosphorus

Plural
uncountable

phosphorus (uncountable)

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Wikipedia has an article on:

Wikipedia

  1. (chemistry) a chemical element (symbol P) with an atomic number of 15, that exists in several allotropic forms.
  2. (obsolete) any substance exhibiting phosphorescence; a phosphor

Usage notes

  • Do not confuse phosphorus with phosphorous, which is an adjective.

Derived terms

Related terms

Translations

See also

External links

For etymology and more information refer to: http://elements.vanderkrogt.net/elem/be.html (A lot of the translations were taken from that site with permission from the author)


Latin

Noun

phosphorus m.

  1. phosphorus

Simple English

File:P,
Red phosphorus in a tube.

[[File:|thumb|White, red, violet, and black phosphorus]] Phosphorus has the chemical symbol P, and its atomic number is 15. Its mass number is 30.97. It is not found in nature as an element but as compounds, such as phosphates. It can be a red or a white waxy solid.

Contents

Properties

Physical properties

Phosphorus comes in several forms. White and red phosphorus are the most common forms. White phosphorus is a waxy white solid. When pure, it is colourless. It is insoluble in water, but soluble in carbon disulphide, an organic solvent. It turns light yellow when in air. It glows in the dark because it "burns" very slowly in air.

When exposed to sunlight, or when heated in its own vapour to 250 °C, it is converted to the red variety. This form does not ignite spontaneously and it is less toxic and less flammable than white phosphorus. The red modification is fairly stable and sublimes with a vapor pressure of 1 atmosphere at 417 °C.

Chemical properties

White phosphorus is more reactive than red phosphorus. White phosphorus catches fire spontaneously in air, burning to make smoke of phosphorus(V) oxide. If it burns in a little air, it produces poisonous phosphorus(III) oxide. When white phosphorus is heated in an alkali, it disproportionates to produce hypophosphites and phosphine. Red phosphorus can burn but needs to be ignited. Phosphorus reacts with the halogens to make phosphorus halides. It reacts with some metals to make phosphides.

Chemical compounds

Phosphorus forms chemical compounds in several oxidation states: -3, +1, +3, and +5. -3 is found in phosphine and phosphides. They are strong reducing agents. +1 compounds are found in hypophosphites. They are strong reducing agents, too. They are toxic. +3 compounds are found in phosphites. They are reducing agents. They are toxic, too. +5 compounds are found in phosphates and phosphoric acid. They are normally non reactive and non toxic. Phosphates are the most common phosphorus compounds. All of the other ones are rare.

Occurrence

It is an essential component of living systems and is found as phosphate in nervous tissue, bones and cell protoplasm. It is also found in the earth as phosphate rock. Phosphate rock is the main source of phosphorus and phosphorus compounds. Many body tissues have calcium phosphates in them.

Preparation

Phosphorus was first made by heating a mixture of phosphates and carbon in an iron pot. The phosphates were made by dissolving bones in strong acids and evaporating the solution.

Phosphorus is made now by heating calcium phosphate, carbon, and silicon dioxide in an electric arc furnace. The heat of the electric arc melts the mixture of materials, and phosphorus gas is given off. It is absorbed under water. This makes white phosphorus.

Uses

As an element

White phosphorus is used in incendiary weapons and smoke grenades. It is also used to make organic compounds that have phosphorus in them. Phosphorus is used to dope semiconductors. Phosphorus is used to remove oxygen from copper. It is also used in making alloys. Red phosphorus is used in matches and flares.

As chemical compounds

Phosphorus compounds are used for fertilizers, soft drinks, toothpaste, and detergents. Most of these are phosphates. Phosphides can be used to kill rodents.

Safety

White phosphorus is very dangerous. It is very toxic and ignites easily, burning with a very hot flame. Red phosphorus is much safer. Some phosphorus compounds are toxic, but the common phosphates are not toxic.

Other pages

frr:Phosphor

Citable sentences

Up to date as of December 18, 2010

Here are sentences from other pages on Phosphorus, which are similar to those in the above article.








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