Phosphorus
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/* As an exception to the octet rule */ minor grammar error, final paragraph in section
:''see [[Phosphorus (morning star)]] for the Greek name of the morning star.''
:''see [[Doctor Phosphorus]] for the Batman villian.''
{{Infobox phosphorus}}
'''Phosphorus''', ({{IPAEng|ˈfɒsfərəs}}), is the [[chemical element]] that has the symbol '''P''' and [[atomic number]] 15. The name comes from the {{lang-el|''φώς''}} (meaning "light") and ''φόρος'' (meaning "bearer"). A [[Valency (chemistry)|multivalent]] [[nonmetal]] of the [[nitrogen group]], phosphorus is commonly found in inorganic [[phosphate minerals|phosphate rocks]].
Due to its high reactivity, phosphorus is never found as a free element in nature on Earth. One form of phosphorus (white phosphorus) emits a faint glow upon exposure to [[oxygen]] — hence its Greek derivation, Φωσφόρος meaning "light-bearer" (Latin ''[[Lucifer]]''), the planet [[Venus]] as "[[Hesperus|Morning Star]]".
Phosphorus is a component of [[DNA]] and [[RNA]] and an essential element for all [[living cell]]s. The most important commercial use of phosphorus-based chemicals is the production of [[fertilizer]]s.
Phosphorus compounds are also widely used in [[explosive material|explosives]], [[nerve agents]], friction [[match]]es, [[firework]]s, [[pesticide]]s, [[toothpaste]], and [[detergent]]s.
== Characteristics==
===Allotropes===
{{main|Allotropes of phosphorus}}
Phosphorus is an excellent example of an element that exhibits [[allotropy]], as its various allotropes have strikingly different properties.
The two most common allotropes are white phosphorus and red phosphorus. A third form, scarlet phosphorus, is obtained by allowing a solution of white phosphorus in carbon disulfide to evaporate in sunlight. A fourth allotrope, black phosphorus, is obtained by heating white phosphorus under very high pressures (12,000 atmospheres). In appearance, properties and structure it is very like graphite, being black and flaky, a conductor of electricity and has puckered sheets of linked atoms. Another allotrope is [[diphosphorus]] - which is highly reactive.
[[Image:White phosphrous molecule.jpg|left|175px]]
[[White phosphorus]] ({{chem|P|4}}) exists as individual molecules made up of four atoms in a [[tetrahedral]] arrangement, resulting in very high [[ring strain]] and instability. It contains 6 single bonds.
White phosphorus is a white, waxy transparent solid. This allotrope 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 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 [[pyrophoricity|pyrophoric]] (self-igniting) upon contact with air as well as [[toxicity|toxic]] (causing severe liver damage on ingestion). The infamous incendiary bomb [[Napalm]] relies, among others, on this principle to spontaneously ignite. 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 P<sub>4</sub>O<sub>10</sub> tetrahedra with oxygen inserted between the phosphorus atoms and at their vertices. White phosphorus is insoluble in water but soluble in carbon disulfide.
The white allotrope can be produced using several different methods. In one process, [[calcium]] phosphate, which is derived from phosphate rock, is heated in an electric or fuel-fired furnace in the presence of [[carbon]] and [[silica]]<ref name="threlfall"/>. 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.
Red phosphorus may be formed by heating white phosphorus to 250°C (482°F) or by exposing white phosphorus to sunlight. Phosphorus after this treatment exists as an [[amorphous]] network of atoms which reduces strain and gives greater stability; further heating results in the red phosphorus becoming crystalline. Red phosphorus does not catch fire in air at temperatures below 240°C, whereas white phosphorus ignites at about 30°C.
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." In addition, a fibrous form exists with similar phosphorus cages. Below is shown a chain of phosphorus atoms which exhibits both the purple and fibrous forms.
[[Image:Hittoff phosphorus chain.jpg|400px]]
One of the forms of red/black phosphorus is a [[cubic]] solid.<ref>{{cite journal
| author = Ahuja, R.
| year = 2003
| title = Calculated high pressure crystal structure transformations for phosphorus
| journal = [[Physica Status Solidi]] Section B
| volume = 235
| issue = 2
| pages = 282–287
| doi = 10.1002/pssb.200301569
}}</ref>
Black phosphorus has an [[Orthorhombic crystal system|orthorhombic]] structure (C<sub>mca</sub>) and is the least reactive allotrope. It consists of many six-membered rings which are interlinked. Each atom is bonded to three other atoms.<ref name="Brown">{{cite journal | author = A. Brown, S. Runquist | journal = [[Acta Crystallogr.]] | volume = 19 | year = 1965 | pages = 684 | doi = 10.1107/S0365110X65004140 | title = Refinement of the crystal structure of black phosphorus}}</ref><ref>{{cite journal | author = Cartz, L.;Srinivasa, S.R.;Riedner, R.J.;Jorgensen, J.D.;Worlton, T.G. | journal = [[Journal of Chemical Physics]] | year = 1979 | volume = 71 | pages = 1718–1721 | doi = 10.1063/1.438523 | title = Effect of pressure on bonding in black phosphorus}}</ref> A recent synthesis of black phosphorus using metal salts as catalysts has been reported.<ref>{{cite journal | author = Stefan Lange, Peer Schmidt, and Tom Nilges | journal = [[Inorg. Chem.]] | year = 2007 | volume = 46 | pages = 4028 | doi = 10.1021/ic062192q | title = Au3SnP7@Black Phosphorus: An Easy Access to Black Phosphorus}}</ref>
The diphosphorus allotrope (P<sub>2</sub>) can be obtained normally only under extreme conditions (for example, from P<sub>4</sub> 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]]).<ref>{{cite journal | journal = [[Science (journal)|Science]] | volume = 313 | issue = 5791 | pages = 1276 | doi = 10.1126/science.1129630 | title = Triple-Bond Reactivity of Diphosphorus Molecules | year = 2006 | author = Piro, N. A. | pmid = 16946068}}</ref>
[[Image:Black phosphorus.jpg|left|200px]]
=== Glow ===
The chemical element [[phosphorus]] ([[Greek language|Greek]]. ''phosphoros'', meaning "light bearer") was discovered by [[Germany|German]] alchemist [[Hennig Brand]] in 1669. 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. However, phosphorus itself is ''not'' a [[phosphor]]; it is highly reactive and gives-off a faint [[chemiluminescence|chemiluminescent]] glow upon uniting with [[oxygen]]. The glow observed by Brand was actually caused by the very slow burning of the phosphorus, but as he saw no flame nor felt any heat he did not recognize it as burning.
The mechanism for that glow was not fully described until 1974.<ref name="shockinghistory">Emsley, John (2000). ''The Shocking History of Phosphorus''. London: Macmillan. ISBN 0-330-39005-8</ref> 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;<ref>[http://nobelprize.org/chemistry/laureates/1956/press.html Nobel Prize in Chemistry 1956 - Presentation Speech], by Professor A. Ölander (committee member)</ref> there is only a range of [[partial pressure]] at which it does. Heat can be applied to drive the reaction at higher pressures.<ref>[http://www.lateralscience.co.uk/phos/index.html Phosphorus Topics page, at Lateral Science]</ref>
In 1974, the glow was explained by R. J. van Zee and A. U. Khan.<ref name="shockinghistory"/> A reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived molecules HPO and P<sub>2</sub>O<sub>2</sub> that both emit visible light. The reaction is slow and only very little of the intermediates is 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 luminescence (glowing by its own reaction, in this case [[chemoluminescence]]), not phosphorescence (re-emitting light that previously fell on it).
=== Isotopes ===<!-- This section is linked from [[Silicon]] -->
{{main|Isotopes of phosphorus}}
{{Expand-section|date=January 2008}}<!-- add about 31P -->
{{splitsection|date=June 2008}}
[[Radioactive decay|Radioactive]] [[isotope]]s of phosphorus include
* <sup>32</sup>P; a [[beta particle|beta]]-emitter (1.71 MeV) with a [[half-life]] of 14.3 days which is used routinely in life-science laboratories, primarily to produce [[radiolabel]]ed DNA and RNA [[probe]]s, ''e.g.'' for use in [[Northern blot]]s or [[Southern blot]]s. Because the high energy beta particles produced penetrate skin and [[cornea]]s, and because any <sup>32</sup>P ingested, inhaled, or absorbed is readily incorporated into bone and [[nucleic acid]]s, [[Occupational Safety and Health Administration]] in the Unites States, and similar institutions in other developed countries require that a [[lab coat]], [[rubber glove|disposable gloves]], and [[safety glasses]] or [[goggle]]s be worn when working with <sup>32</sup>P, 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, [[radiation shield|shielding]] this radiation with the normally used dense materials (''e.g.'' [[lead]]), gives rise to secondary emission of [[X-ray]]s 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]].<ref>http://www.oseh.umich.edu/TrainP32.pdf</ref>
* <sup>33</sup>P; 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|Phosphate minerals]].''
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 [[mineral]]s.
Phosphate rock, which is partially made of [[apatite]] (an impure tri-calcium phosphate mineral), is an important commercial source of this element. About 50 per cent of the global phosphorus reserves are in the Arab nations.<ref>http://www.anba.com.br/ingles/noticia.php?id=17288</ref> 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<ref name="threlfall"/>. In the early 1990s Albright and Wilson's purified wet phosphoric acid business was being affected by phosphate rock sales by China and the entry of their long standing Moroccan phosphate suppliers into the purified wet phosphoric acid business.<ref name="podger">Podger, Hugh, (2002). ''Albright & Wilson: The Last 50 Years''. Studley: Brewin Books. ISBN 1-85858-223-7</ref>
In 2007, at the then current rate of consumption, the supply of phosphorus was estimated to run out in 345 years.<ref>{{cite journal | date=[[May 26]], [[2007]] | journal = New Scientist | volume = 194 | issue = 2605 | pages = 38–39 | issn = 0262 4079 | title = How Long Will it Last? }}</ref> 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."<ref>{{cite news | url = http://business.timesonline.co.uk/tol/business/industry_sectors/natural_resources/article4193017.ece | title = Scientists warn of lack of vital phosphorus as biofuels raise demand | date = 2008-06-23 | author = Leo Lewis | publisher = [[The Times]]}}</ref>
== Compounds ==
{{Expand-section|date=January 2008}}
''See also [[:Category:Phosphorus compounds|Phosphorus compounds]]''
<div style="-moz-column-count:3; column-count:3;">
* [[Hydride]]: [[Phosphine|PH<sub>3</sub>]]
* [[Halide]]s: [[Phosphorus pentabromide|PBr<sub>5</sub>]], [[Phosphorus tribromide|PBr<sub>3</sub>]], [[Phosphorus trichloride|PCl<sub>3</sub>]], [[Phosphorus triiodide|PI<sub>3</sub>]]
* [[Oxide]]s:[[Phosphorus trioxide|P<sub>4</sub>O<sub>6</sub>]], [[Phosphorus pentoxide|P<sub>4</sub>O<sub>10</sub>]]
* [[Sulfide]]s: [[Phosphorus pentasulfide|P<sub>2</sub>S<sub>5</sub>]], [[Phosphorus sesquisulfide|P<sub>4</sub>S<sub>3</sub>]]
* [[Acid]]s: [[Hypophosphorous acid|H<sub>3</sub>PO<sub>2</sub>]], [[Phosphoric acid|H<sub>3</sub>PO<sub>4</sub>]]
* [[Phosphate]]s: [[Ammonium phosphate|(NH<sub>4</sub>)<sub>3</sub>PO<sub>4</sub>]], [[Tricalcium phosphate|Ca<sub>3</sub>(PO<sub>4</sub>)<sub>2</sub>)]], [[Iron(III) phosphate|FePO<sub>4</sub>]], [[Iron(II) phosphate|Fe<sub>3</sub>(PO<sub>4</sub>)<sub>2</sub>]], [[Trisodium phosphate|Na<sub>3</sub>PO<sub>4</sub>]], [[Calcium dihydrogen phosphate|Ca(H<sub>2</sub>PO<sub>4</sub>)<sub>2</sub>]], [[Monopotassium phosphate|KH<sub>2</sub>PO<sub>4</sub>]]
* [[Phosphide]]s: [[Calcium phosphide|Ca<sub>3</sub>P<sub>2</sub>]], [[Gallium(III) phosphide|GaP]], [[Zinc phosphide|Zn<sub>3</sub>P<sub>2</sub>]] [[Copper phosphide|Cu<sub>3</sub>P]]
* [[Organophosphorus]] and [[organophosphate]]s: [[Lawesson's reagent]], [[Parathion]], [[Sarin]], [[Soman]], [[Tabun (nerve agent)|Tabun]], [[Triphenyl phosphine]], [[VX (nerve agent)|VX]] nerve gas
</div>
=== As an exception to the octet rule ===
{{Details|Octet rule}}
The simple [[Lewis structure]] for the [[trigonal bipyramid molecular geometry|trigonal bipyramidal]] [[phosphorus pentachloride|PCl<sub>5</sub>]] 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 bond|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 P.
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 (the majority, in fact) to which it just does not apply.
== Applications ==
{{Wikify|date=January 2008}}
* Phosphorous being an essential plant nutrient, phosphorus 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% P<sub>2</sub>O<sub>5</sub>. Global demand for fertilizers led to large increase in [[phosphate]] (PO<sub>4</sub><sup>3-</sup>) production in the second half of the 20th century.
* Phosphorus is widely used to make [[organophosphorus compound]]s, through the intermediates [[phosphorus chlorides]] and the two phosphorus sulfides: [[phosphorus pentasulfide]], and [[phosphorus sesquisulfide]].<ref name="threlfall"/> Organophosphorus compounds have many applications, including in [[plasticizers]], [[flame retardant]]s, [[pesticide]]s, [[extraction agent]]s, and water treatment.
* Phosphorus is also an important component in [[steel]] production, in the making of [[phosphor bronze]], and in many other related products.
* Phosphates are utilized in the making of special [[glass]]es that are used for [[sodium lamp]]s.
* Bone-ash, [[calcium phosphate]], is used in the production of fine china.
* [[Sodium tripolyphosphate]] made from phosphoric acid is used in laundry detergents in some countries, but banned for this use in others.
* Phosphoric acid made from elemental phosphorus is used in food applications such as some soda beverages. The acid is also a starting point to make food grade phosphates.<ref name="threlfall">Threlfall, R.E., (1951). ''100 years of Phosphorus Making: 1851 - 1951''. Oldbury: [[Albright and Wilson]] Ltd</ref> These include mono-calcium phosphate which is employed in [[baking powder]] and [[sodium tripolyphosphate]] and other sodium phosphates<ref name="threlfall"/>. Among other uses these are used to improve the characteristics of processed meat and cheese. Others are used in toothpaste.<ref name="threlfall"/> [[Trisodium phosphate]] is used in cleaning agents to [[water softener|soften water]] and for preventing pipe/boiler tube [[corrosion]].
* [[white phosphorus incendiary|White phosphorus]] is used in [[military]] applications as [[incendiary device|incendiary bomb]]s, for [[smoke-screen]]ing as smoke pots and [[smoke bomb]]s, and in [[tracer ammunition]].
* Red phosphorus is essential for manufacturing matchbook strikers, flares,<ref name="threlfall"/> safety matches, pharmaceutical grade and street [[methamphetamine]], and is used in [[cap gun]] caps.
* Phosphorus sesquisulfide is used in heads of strike-anywhere matches.<ref name="threlfall"/>
* In trace amounts, phosphorus is used as a [[dopant]] for [[N-type semiconductor]]s.
* <sup>32</sup>P and <sup>33</sup>P are used as radioactive tracers in biochemical laboratories (see [[#Isotopes|Isotopes]]).
== Biological role ==
Phosphorus is a key element in all known forms of [[life]]. Inorganic phosphorus in the form of the phosphate PO<sub>4</sub><sup>3-</sup> 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. [[Phospholipid]]s are the main structural components of all cellular membranes. [[Calcium phosphate]] salts assist in stiffening [[bone]]s.
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 in the form of phosphate. Only about 0.1% of body phosphate circulates in the blood, but this amount reflects the amount of phosphate available to soft tissue cells.
In medicine, low phosphate syndromes are caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes which draw phosphate from the blood or pass too much of it into the urine. All are characterized by [[hypophosphatemia]] (see article for medical details). Symptoms of low phosphate include muscle and neurological dysfunction, and disruption of muscle and blood cells due to lack of ATP.
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]].
== History ==
Phosphorus ([[Greek language|Greek]] ''phosphoros'' was the ancient name for the planet [[Venus (planet)|Venus]], but in [[Greek mythology]], Hesperus and Eosphorus could be confused with Phosphorus) was discovered by German [[alchemy|alchemist]] [[Hennig Brand]] in 1669 through a preparation from [[urine]], which contains considerable quantities of dissolved phosphates from normal metabolism. Working in [[Hamburg]], Brand attempted to create the infamous [[Philosopher's stone]] through the [[distillation]] of some [[salt]]s by evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. Since that time, phosphorescence has been used to describe substances that shine in the dark without burning.
Phosphorus was recognized as a chemical element at the emergence of the [[Atomic_theory#Modern_atomic_theory|atomic theory]] that gradually occurred in the late part of the 18th century and the early 19th century, and was formulated by [[John Dalton]].
Phosphorus was first made commercially, for the match industry, in the 19th century, by distilling off phosphorus vapor from precipitated phosphates heated in a [[retort]].<ref name="threlfall"/> The precipitated phosphates were made from ground-up bones that had been de-greased and treated with strong acids.<ref name="threlfall"/> This process became obsolete in the late 1890s when the [[electric arc furnace]] was adapted to reduce phosphate rock.<ref name="threlfall"/>
Early matches used white phosphorus in their composition, which was dangerous due to its toxicity. Murders, suicides and accidental [[poison]]ings 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).<ref name="shockinghistory"/> 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 a [[Berne]] Convention, requiring its adoption as a safer alternative for match manufacture.
The electric furnace method allowed production to increase to the point where phosphorus could be used in weapons of war.<ref name="shockinghistory"/><ref name="threlfall"/> In [[World War I]] it was used in incendiaries, [[smoke screen]]s and tracer bullets.<ref name="threlfall"/> A special incendiary bullet was developed to shoot at [[hydrogen]]-filled [[Zeppelin]]s over [[United Kingdom|Britain]] (hydrogen being highly [[inflammable]] if it can be ignited).<ref name="threlfall"/> During [[World War II]], [[Molotov cocktail]]s 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|precautions]] below). People covered in it have been known to commit suicide due to the torment.
Today phosphorus production is larger than ever. It is used as a precursor for various chemicals,<ref>{{cite journal|title=The American Phosphorus Industry|author =Aall C. H.|journal =Industrial & Engineering Chemistry|year= 1952|volume =44|doi =10.1021/ie50511a018|issue=7|pages=1520–1525}}</ref> 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]].
=== 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 P<sup>3+</sup> valency: so, just as [[sulfur]] forms sulfur'''ous''' and sulfur'''ic''' compounds, phosphor'''us''' forms phosphor'''ous''' compounds (see e.g. [[phosphorous acid]]) and P<sup>5+</sup> valency phosphor'''ic''' compounds (see e.g. [[Phosphoric acids and phosphates]]).
== Precautions ==
[[Image:Hazard F.svg|left|100px]]
[[Image:Skull and crossbones.svg|right|80px]]
Organic compounds of phosphorus form a wide class of materials, some of which are extremely toxic. [[Fluorophosphate]] [[ester]]s are among the most potent [[neurotoxin]]s known. A wide range of organophosphorus compounds are used for their toxicity to certain organisms as [[pesticides]] ([[herbicides]], [[insecticides]], [[fungicides]], etc.) and [[weapon]]ized as nerve agents. Most inorganic phosphates are relatively nontoxic and essential nutrients. For environmentally adverse effects of phosphates see [[eutrophication]] and [[algal bloom]]s.
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". <ref>[http://www.emedicine.com/EMERG/topic918.htm emedicine.com] CBRNE - Incendiary Agents, White Phosphorus (Smoking Stool Syndrome)</ref>
When the white form is exposed to sunlight or when it is heated in its own vapour to 250°C, it is transmuted to the red form, which does not phosphoresce 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 [[oxide]]s when it is heated.
Upon exposure to elemental phosphorus, in the past it was suggested to wash the affected area with 2% [[copper sulfate]] solution to form harmless compounds that can be washed away. According to the recent ''US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries'', "Cupric (copper(II)) sulfate has been used by U.S. personnel in the past and is still being used by some nations. However, copper sulfate is toxic and its use will be discontinued. Copper sulfate may produce kidney and cerebral toxicity as well as intravascular hemolysis."<ref>[http://www.vnh.org/FM8285/Chapter/chapter9.html US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries]</ref>
[[Image:Phosphorus explosion.gif|thumb|Phosphorus explosion]]
The manual suggests instead "a bicarbonate solution to neutralize phosphoric acid, which will then allow removal of visible WP. Particles often can be located by their emission of smoke when air strikes them, or by their phosphorescence in the dark. In dark surroundings, fragments are seen as luminescent spots." Then, "Promptly debride the burn if the patient's condition will permit removal of bits of WP which might be absorbed later and possibly produce systemic poisoning. DO NOT apply oily-based ointments until it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns." As white phosphorus readily mixes with oils, any oily substances or ointments are not recommended until the area is thoroughly cleaned and all white phosphorus removed.
Further warnings of toxic effects and recommendations for treatment can be found in the ''Emergency War Surgery NATO Handbook: Part I: Types of Wounds and Injuries: Chapter III: Burn Injury: Chemical Burns And White Phosphorus injury''.<ref>[http://www.brooksidepress.org/Products/OperationalMedicine/DATA/operationalmed/Manuals/NATOEWS/ch03/03ChemicalBurns.html Emergency War Surgery NATO Handbook: Part I: Types of Wounds and Injuries: Chapter III: Burn Injury: Chemical Burns And White Phosphorus injury].</ref>
=== 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.<ref>{{cite journal | author = Skinner, H.F. | year = 1990 | title = Metamphetamine synthesis via hydriodic acid/red phosphorus reduction of ephedrine | journal = [[Forensic Science International]] | volume = 48 | issue = 2 | pages = 123–134 | doi = 10.1016/0379-0738(90)90104-7}}</ref> For this reason, two allotropes of elemental phosphorus—red phosphorus and white phosphorus—were designated by the [[United States of America|United States]] [[Drug Enforcement Administration]] as [[DEA list of chemicals|List I precursor chemicals]] under [[Code of Federal Regulations|21 CFR 1310.02]] effective November 17, 2001.<ref name="66 CFR 52670">[http://frwebgate.access.gpo.gov/cgi-bin/getdoc.cgi?dbname=2001_register&docid=01-26013-filed 66 FR 52670—52675.] 17 October 2001.</ref> 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.<ref name="66 CFR 52670"/><ref name="21 CFR 1309">[http://www.access.gpo.gov/nara/cfr/waisidx_06/21cfr1309_06.html 21 CFR 1309]</ref><ref name="CSA">[http://www.usdoj.gov/dea/pubs/csa.html 21 USC, Chapter 13 (Controlled Substances Act)]</ref>
== References ==
{{reflist}}
== External links ==
{{commonscat}}
* [http://periodic.lanl.gov/elements/15.html Los Alamos National Laboratory – Phosphorus]
* [http://www.webelements.com/webelements/elements/text/P/index.html WebElements.com: Phosphorus]
* [http://www.ncbi.nlm.nih.gov/entrez/query.fcgi?cmd=Retrieve&db=PubMed&list_uids=946251&dopt=Abstract; Entrez PubMed: Acute Yellow Phosphorus Poisoning]
* [http://www.emedicine.com/EMERG/topic918.htm eMedicine.com: Article on White Phophorus as used as weapon]
* [http://www.phosphorus-recovery.tu-darmstadt.de Website of the Technische Universität Darmstadt and the CEEP about Phosphorus Recovery]
{{Clear}}
{{Compact periodic table}}
[[Category:Chemical elements]]
[[Category:DEA List I chemicals]]
[[Category:Dietary minerals]]
[[Category:Nonmetals]]
[[Category:Pnictogens]]
[[Category:Phosphorus|*]]
[[Category:Pyrotechnic fuels]]
[[af:Fosfor]]
[[ar:فوسفور]]
[[ast:Fósforu]]
[[az:Fosfor]]
[[bn:ফসফরাস]]
[[be:Фосфар]]
[[bs:Fosfor]]
[[bg:Фосфор]]
[[ca:Fòsfor]]
[[cs:Fosfor]]
[[co:Fosfaru]]
[[cy:Ffosfforws]]
[[da:Fosfor]]
[[de:Phosphor]]
[[dv:ފޮސްފަރަސް]]
[[et:Fosfor]]
[[el:Φωσφόρος]]
[[es:Fósforo]]
[[eo:Fosforo]]
[[eu:Fosforo]]
[[fa:فسفر]]
[[fr:Phosphore]]
[[fur:Fosfar]]
[[ga:Fosfar]]
[[gv:Fosfaar]]
[[gl:Fósforo (elemento)]]
[[ko:인]]
[[hy:Ֆոսֆոր]]
[[hi:फास्फोरस]]
[[hr:Fosfor]]
[[io:Fosfo]]
[[id:Fosfor]]
[[is:Fosfór]]
[[it:Fosforo]]
[[he:זרחן]]
[[sw:Posferi]]
[[ht:Fosfò]]
[[ku:Fosfor]]
[[la:Phosphorus]]
[[lv:Fosfors]]
[[lb:Phosphor]]
[[lt:Fosforas]]
[[jbo:sackycmu]]
[[hu:Foszfor]]
[[mk:Фосфор]]
[[ml:ഫോസ്ഫറസ്]]
[[mi:Pūtūtae-whetū]]
[[mr:फॉस्फरस]]
[[ms:Fosforus]]
[[nl:Fosfor]]
[[ja:リン]]
[[no:Fosfor]]
[[nn:Fosfor]]
[[nov:Fosfore]]
[[oc:Fosfòr]]
[[uz:Fosfor]]
[[nds:Phosphor]]
[[pl:Fosfor]]
[[pt:Fósforo]]
[[ro:Fosfor]]
[[qu:Phusphuru]]
[[ru:Фосфор]]
[[sq:Fosfori]]
[[scn:Fosfuru]]
[[simple:Phosphorus]]
[[sk:Fosfor]]
[[sl:Fosfor]]
[[sr:Фосфор]]
[[sh:Fosfor]]
[[fi:Fosfori]]
[[sv:Fosfor]]
[[ta:பாஸ்பரஸ்]]
[[te:భాస్వరము]]
[[th:ฟอสฟอรัส]]
[[vi:Phốtpho]]
[[tg:Фосфор]]
[[tr:Fosfor]]
[[uk:Фосфор]]
[[ur:فاسفورس]]
[[zh:磷]]