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{{About|the chemical element}}
{{Use British English|date=January 2018}}
{{Infobox phosphorus|engvar=en-GB}}

'''Phosphorus''' is a [[chemical element]]; it has [[Chemical symbol|symbol]] '''P''' and [[atomic number]] 15. All elemental forms of phosphorus are highly [[Reactivity (chemistry)|reactive]] and are therefore never found in nature. They can nevertheless be prepared artificially, the two most common [[allotropes]] being [[white phosphorus]] and [[red phosphorus]]. With {{chem2|^{31}P}} as its only stable [[isotope]], phosphorus has an occurrence in [[Earth's crust]] of about 0.1%, generally as [[phosphate rock]]. A member of the [[pnictogen]] family, phosphorus readily forms a wide variety of [[organic compound|organic]] and [[inorganic compound|inorganic]] compounds, with as its main [[oxidation state]]s +5, +3 and −3.

The isolation of white phosphorus in 1669 by [[Hennig Brand]] marked the scientific community's first discovery since Antiquity of an element. The name phosphorus is a reference to the [[Phosphorus (morning star)|god of the Morning star]] in [[Greek mythology]], inspired by the faint glow of white phosphorus when exposed to [[oxygen]]. This property is also at the origin of the term ''[[phosphorescence]]'', meaning glow after illumination, although white phosphorus itself does not exhibit phosphorescence, but [[chemiluminescence]] caused by its [[oxidation]]. Its high [[toxicity]] makes exposure to white phosphorus very dangerous, while its [[flammability]] and [[pyrophoricity]] can be weaponised in the form of [[incendiaries]]. Red phosphorus is less dangerous and is used in [[match]]es and [[fire retardant]]s.

Most industrial production of phosphorus is focused on the mining and transformation of phosphate rock into [[phosphoric acid]] for [[phosphate]]-based [[fertiliser]]s. Phosphorus is an essential and often [[limiting nutrient]] for plants, and while natural levels are normally maintained over time by the [[phosphorus cycle]], it is too slow for the regeneration of soil that undergoes [[intensive cultivation]]. As a consequence, these fertilisers are vital to modern agriculture. The leading producers of phosphate ore in 2024 were China, Morocco, the United States and Russia, with two-thirds of the estimated exploitable phosphate reserves worldwide in Morocco alone. Other applications of phosphorus compounds include [[pesticide]]s, [[food additive]]s, and [[detergent]]s.

Phosphorus is essential to all known forms of [[life]], largely through [[organophosphate]]s, organic compounds containing the phosphate ion {{chem2|PO4(3−)}} as a [[functional group]]. These include [[DNA]], [[RNA]], [[Adenosine triphosphate|ATP]], and [[phospholipid]]s, complex compounds fundamental to the functioning of all [[Cell (biology)|cells]]. The main component of bones and teeth, [[bone mineral]], is a modified form of [[hydroxyapatite]], itself a phosphorus mineral.

==History==
[[File:Joseph Wright of Derby The Alchemist.jpg|thumb|upright|left|''[[The Alchemist Discovering Phosphorus|The Alchemist in Search of the Philosophers Stone]]'' (1771), by [[Joseph Wright of Derby|Joseph Wright]], depicting Hennig Brand discovering phosphorus.]]
Phosphorus was the [[Discovery of chemical elements|first element to be "discovered"]], in the sense that it was not known since ancient times.<ref name="Weeks1932" /> The discovery is credited to the [[Hamburg]] alchemist [[Hennig Brand]] in 1669, who was attempting to create the fabled [[philosopher's stone]].<ref name="Beatty2000" /> To this end, he experimented with [[urine]], which contains considerable quantities of dissolved phosphates from normal metabolism.<ref name="Mellor1939" /> By letting the urine rot (a step later discovered to be unnecessary),<ref name="Sommers2007" /> boiling it down to a paste, then [[distillation|distilling]] it at a high temperature and leading the resulting vapours through water, he obtained a white, waxy substance that glowed in the dark and burned brilliantly. He named it in {{langx|la|phosphorus mirabilis|lit=miraculous bearer of light}}. The word phosphorus itself ({{langx|grc|Φωσφόρος|Phōsphoros|lit=light-bearer}}) originates from [[Greek mythology]], where it references the [[Phosphorus (morning star)|god of the morning star]], also known as the planet [[Venus]].<ref name="Mellor1939" /><ref name="Schmundt2010" />

Brand at first tried to keep the method secret,<ref name="Stillman1960" /> but later sold the recipe for 200 [[thaler]]s to {{ill|Johann Daniel Kraft|de}} from [[Dresden]].<ref name="Mellor1939" /> Kraft toured much of Europe with it, including [[London]], where he met with [[Robert Boyle]]. The crucial fact that the substance was made from urine was eventually found out, and [[Johann von Löwenstern-Kunckel|Johann Kunckel]] was able to reproduce it in Sweden in 1678. In 1680, Boyle also managed to make phosphorus and published the method of its manufacture.<ref name="Mellor1939" /> He was the first to use phosphorus to ignite [[sulfur]]-tipped wooden splints, forerunners of modern matches,<ref name="Baccini2012" /> and also improved the process by using sand in the reaction:
:{{chem2|4 NaPO3 + 2 SiO2 + 10 C -> 2 Na2SiO3 + 10 CO + P4}}
Boyle's assistant [[Ambrose Godfrey|Ambrose Godfrey-Hanckwitz]] later made a business of the manufacture of phosphorus.

In 1777, [[Antoine Lavoisier]] recognised phosphorus as an element after [[Johan Gottlieb Gahn]] and [[Carl Wilhelm Scheele]] showed in 1769 that [[calcium phosphate]] is found in bones by obtaining elemental phosphorus from [[bone ash]].<ref name="LavMem1rySrc" /> Bone ash subsequently became the primary industrial source of phosphorus and remained so until the 1840s.<ref name="Wagner1897" /> The process consisted of several steps.<ref name="Thomson1870" /><ref name="Threlfall1951" /> First, grinding up the bones into their constituent [[tricalcium phosphate]] and treating it with [[sulfuric acid]]:
:{{chem2|Ca3(PO4)2 + 2 H2SO4 -> Ca(H2PO4)2 + 2 CaSO4}}
Then, dehydrating the resulting [[monocalcium phosphate]]:
:{{chem2|Ca(H2PO4)2 -> Ca(PO3)2 + 2 H2O}}
Finally, mixing the obtained calcium [[metaphosphate]] with ground [[coal]] or [[charcoal]] in an iron pot, and distilling phosphorus vapour out of a [[retort]]:
:{{chem2|3 Ca(PO3)2 + 10 C -> Ca3(PO4)2 + 10 CO + P4}}
This way, two-thirds of the phosphorus was turned into white phosphorus while one-third remained in the residue as calcium [[orthophosphate]]. The [[carbon monoxide]] produced during the reaction process was burnt off in a [[Gas flare|flare stack]].

[[File:DSCN5766-guano-glantz crop b.jpg|thumb|upright|left|Guano mining in the Central [[Chincha Islands]], c. 1860]]
[[File:Phosphorus bottle pocket matches, 1828 - Joseph Allen Skinner Museum - DSC07746.JPG|thumb|left|Matches from 1828. The sulfur-tipped match is dipped into liquid containing white phosphorus, and ignites as it is pulled out of the bottle.]]
[[File:Phosphate smelting furnace worker, Muscle Shoals fsac.1a35278u.jpg|thumb|left|A worker tends an electric phosphate smelting furnace in [[Muscle Shoals, Alabama]], 1942]]
[[File:Phosphorus explosion.gif|thumb|left|White phosphorus shell explosion in France during the First World War (1918)]]

In 1609 [[Inca Garcilaso de la Vega]] wrote the book ''Comentarios Reales'' in which he described many of the agricultural practices of the Incas prior to the arrival of the Spaniards and introduced the use of [[guano]] as a [[fertiliser]]. As Garcilaso described, the Incas near the coast harvested guano.<ref name="Leigh2004" /> In the early 1800s [[Alexander von Humboldt]] introduced guano as a source of agricultural fertiliser to Europe after having discovered it in exploitable quantities on islands off the coast of [[South America]]. It has been reported that, at the time of its discovery, the guano on some islands was over 30&nbsp;meters deep.<ref name="Skaggs1995" /> The guano had previously been used by the [[Moche culture|Moche]] people as a source of fertiliser by mining it and transporting it back to [[Peru]] by boat. International commerce in guano did not start until after 1840.<ref name="Skaggs1995" /> By the start of the 20th century guano had been nearly completely depleted and was eventually overtaken with the discovery of methods of production of [[superphosphate]].

[[Match#History|Early matches]] used white phosphorus in their composition, and were very dangerous due to both its toxicity and the way the match was ignited. The first striking match with a phosphorus head was invented by [[Charles Sauria]] in 1830. These matches (and subsequent modifications) were made with heads of white phosphorus, an oxygen-releasing compound ([[potassium chlorate]], [[lead dioxide]], or sometimes [[nitrate]]), and a binder. They were poisonous to the workers in manufacture, exposure to the vapours causing  severe [[necrosis]] of the bones of the jaw, known as "[[phossy jaw]]".<ref name="Hughes1962" /> Additionally, they were sensitive to storage conditions, toxic if ingested, and hazardous when accidentally ignited on a rough surface.<ref name="Crass1941" /><ref name="Oliver1996" /> The very high risks for match workers was at the source of several notable early cases of [[industrial action]], such as the 1888 London [[Matchgirls' strike]].

The discovery of red phosphorus allowed for the development of matches that were both much safer to use and to manufacture, leading to the gradual replacement of white phoshphorus in matches. Additionally, around 1900 French chemists Henri Sévène and Emile David Cahen invented the modern strike-anywhere match, wherein the white phosphorus was replaced by phosphorus sesquisulfide ({{chem2|P4S3}}), a non-toxic and non-pyrophoric compound that ignites under friction. For a time these safer strike-anywhere matches were quite popular but in the long run they were superseded by the modern red phosphorus-based safety match. Following the implementation of these new manufacturing methods, production of white phosphorus matches was banned in several countries between 1872 and 1925,<ref name="Charnovitz1987" /> and an international [[treaty]] to this effect was signed following the [[Berne Convention (1906)]].<ref name="Goldfrank2006" />

[[Phosphate rock]], which usually contains calcium phosphate, was first used in 1850 to make phosphorus. With the introduction of the [[submerged-arc furnace for phosphorus production]] by [[James Burgess Readman]] in 1888<ref name="Toy1975" /> (patented 1889),<ref name="Patent" /> the use of bone-ash became obsolete.<ref name="Threlfall1951" /><ref name="Mellor1939" /> After the depletion of world guano sources about the same time, mineral phosphates became the major source of phosphate fertiliser production. Phosphate rock production greatly increased after World War II, and remains the primary global source of phosphorus and phosphorus chemicals today.

The electric furnace method allowed production to increase to the point where it became possible that [[White phosphorus munition#History|white phosphorus could be weaponised in war]]. In [[World War I]], it was used in [[incendiary ammunition]], [[smoke screen]]s and [[tracer ammunition]]. A special incendiary bullet was developed to shoot at [[hydrogen]]-filled [[Zeppelin]]s over Britain (hydrogen being highly [[flammable]]).<ref name="Threlfall1951" />

During [[World War II]], [[Molotov cocktail]]s made of phosphorus dissolved in [[petrol]] were distributed in Britain to specially selected civilians as part of the [[British anti-invasion preparations of the Second World War|preparations for a potential invasion]]. The United States also developed the M15 white-phosphorus hand grenade, a precursor to the [[M34 grenade]], while the British introduced the similar [[No 77 grenade]]. These multipurpose grenades were mostly used for signaling and smoke screens, although they were also efficient [[anti-personnel weapon]]s.<ref name="Dockery1997" /> The difficulty of extinguishing burning phosphorus and the very severe burns it causes had a strong psychological impact on the enemy.<ref name="Greenwood1997" /> Phosphorus [[incendiary bomb]]s were used on a large scale, notably to [[Bombing of Hamburg in World War II|destroy Hamburg]], the place where the "miraculous bearer of light" was first discovered.<ref name="Schmundt2010" />

==Characteristics==
===Isotopes===
{{Main|Isotopes of phosphorus}}
There are 22 known [[isotope]]s of phosphorus,{{NUBASE2016|ref}} ranging from {{chem2|^{26}P}} to {{chem2|^{47}P}}.<ref name="Neufcourt2019" /> Only {{chem2|^{31}P}} is stable and is therefore present at 100% abundance. The half-integer [[nuclear spin]] and high abundance of {{chem2|^{31}P}} make [[phosphorus-31 nuclear magnetic resonance]] spectroscopy a very useful analytical tool in studies of phosphorus-containing samples.

Two [[radioactive isotope]]s of phosphorus have half-lives suitable for biological scientific experiments, and are used as radioactive tracers in biochemical laboratories.<ref name="Atwood2013" /> These are:
* {{chem2|^{32}P|link=phosphorus-32}}, a [[beta particle|beta]]-emitter (1.71&nbsp;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 [[Hybridization probe|probes]], e.g. for use in [[Northern blot]]s or [[Southern blot]]s. 
* {{chem2|^{33}P}}, a beta-emitter (0.25&nbsp;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.
The high-energy beta particles from {{chem2|^{32}P}} penetrate skin and [[cornea]]s and any {{chem2|^{32}P}} ingested, inhaled, or absorbed is readily incorporated into bone and [[nucleic acid]]s. For these reasons, personnel working with {{chem2|^{32}P}} is required to wear lab coats, disposable gloves, and safety glasses, and avoid working directly over open containers. [[Biomonitoring|Monitoring]] personal, clothing, and surface contamination is also required. The high energy of the beta particles gives rise to secondary emission of [[X-ray]]s via [[Bremsstrahlung]] (braking radiation) in dense shielding materials such as lead. Therefore, the radiation must be [[Radiation protection|shielded]] with low density materials such as water, acrylic or other plastic.<ref name="OSEH" />

===Atomic properties===
A phosphorus atom has 15 electrons, 5 of which are [[valence electron]]s. This results in the [[electron configuration]] 1s<sup>2</sup>2s<sup>2</sup>2p<sup>6</sup>3s<sup>2</sup>3p<sup>3</sup>, often simplified as [Ne]3s<sup>2</sup>3p<sup>3</sup>, omitting the [[core electron]]s which have a configuration equivalent to the [[noble gas]] of the preceding [[period (periodic table)|period]], in this case [[neon]]. The molar [[ionisation energies]] of these five electrons are 1011.8, 1907, 2914.1, 4963.6 and 6273.9&nbsp;k[[joule per mole|J⋅mol<sup>−1</sup>]].

Phosphorus is a member of the [[pnictogen]]s (also called [[group (periodic table)|group]] 15) and [[period 3 element]]s, and many of its chemical properties can be inferred from its position on the [[periodic table]] as a result of [[periodic trends]]. Like [[nitrogen]], [[arsenic]] and [[antimony]], its main [[oxidation state]]s are −3, +3 and +5, with every one in-between less common but known. Phosphorus shows as expected more [[electronegativity]] than [[silicon]] and arsenic, less than [[sulfur]] and nitrogen, but also notably less than [[carbon]], affecting the nature and properties of P–C bonds. It is the element with the lowest [[atomic number]] to exhibit [[hypervalence]], meaning that it can form more [[chemical bond|bond]]s per atom that would normally be permitted by the [[octet rule]].

===Allotropes===
{{Main|Allotropes of phosphorus}}
{{multiple image|perrow=2|total_width=320|caption_align=center
| header = {{font|size=100%|font=Sans-serif|text=Crystalline structures of the main phosphorus allotropes}}
| align = right
| image_style = border:none;
|image1=White phosphorus molecule.jpg
 |alt1=
 |caption1=White
|image2=redPhosphorus.jpg
 |alt2=
 |caption2=Red
|image3=Hittorf's violet phosphorus.png
 |alt3=
 |caption3=Violet
|image4=BlackPhosphorus.jpg
 |alt4=
 |caption4=Black
}}
Phosphorus has several [[allotropy|allotropes]] that exhibit very diverse properties.<ref name="Holleman1985" /> The most useful and therefore common is [[white phosphorus]], followed by [[red phosphorus]]. The two other main allotropes, violet and black phosphorus, have either a more fundamental interest or specialised applications. Many other allotropes have been theorised and synthesised, with the search for new materials an active area of research.<ref name="Tian2023" /> Commonly mentioned "yellow phosphorus" is not an allotrope, but a result of the gradual degradation of white phosphorus into red phosphorus, accelerated by light and heat. This causes white phosphorus that is aged or otherwise impure (e.g. weapons-grade) to appear yellow.

White phosphorus is a soft, waxy [[molecular solid]] that is insoluble in water.<ref name="Greenwood1997" /> It is also very toxic, highly [[flammable]] and [[pyrophoricity|pyrophoric]], igniting in air at about {{convert|30|C|K}}.<ref name="Mellor1939" /> Structurally, it is composed of {{chem2|P4}} [[tetrahedra]]. The nature of bonding in a given {{chem2|P4}} tetrahedron can be described by [[spherical aromaticity]] or cluster bonding, that is the electrons are highly [[Delocalized electron|delocalized]]. This has been illustrated by calculations of the magnetically induced currents, which sum up to 29&nbsp;nA/T, much more than in the archetypical [[Aromaticity|aromatic]] molecule [[benzene]] (11&nbsp;nA/T).<ref name="Cossairt2010" /> The {{chem2|P4}} molecule in the gas phase has a P-P bond length of 2.1994(3)&nbsp;Å as determined by [[gas electron diffraction]].<ref name="Cossairt2010" /> White phosphorus exists in two crystalline forms named α (alpha) and β (beta), differing in terms of the relative orientation of the constituent {{chem2|P4}} tetrahedra.<ref name="Roberts1992" /><ref name="Averbuch-Pouchot1996" /> The α-form is most stable at room temperature and has a [[cubic crystal structure]]. When cooled down to {{convert|195.2|K|C}} it transforms into the β-form, turning into an [[hexagonal crystal structure]]. When heated up, the tetrahedral structure is conserved after melting at {{convert|317.3|K|C}} and boiling at {{convert|553.7|K|C}}, before facing [[thermal decomposition]] at {{convert|1100|K|C}} where it turns into gaseous [[diphosphorus]] ({{chem2|P2}}).<ref name="Arndt1997" /> This molecule contains a triple bond and is analogous to {{chem2|N2}}; it can also be generated as a transient intermediate in solution by thermolysis of organophosphorus precursor reagents.<ref name="Piro2006" /> At still higher temperatures, {{chem2|P2}} dissociates into atomic P.<ref name="Greenwood1997" />

[[File:White phosphorus glowing e17.png|thumb|right|White phosphorus exposed to air glows in the dark.]]
When exposed to air, white phosphorus faintly glows green and blue due to [[oxidation]], a phenomenon best visible in the dark. This reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived molecules {{chem2|HPO}} and {{chem2|P2O2}} that both emit visible light.<ref name="Vanzee1976" /> However, in a pure-oxygen environment phosphorus does not glow at all, with the oxidation happening only in a range of [[partial pressure]]s.<ref name="Ölander1956" /> Derived from this phenomenon, the terms ''[[phosphor]]s'' and ''[[phosphorescence]]'' have been loosely used to describe substances that shine in the dark. However, phosphorus itself is not phosphorescent but [[chemiluminescent]], since it glows due to a chemical reaction and not the progressive reemission of previously absorbed light.<ref name="Sommers2007" />

Red phosphorus is [[polymer]]ic in structure. It can be viewed as a derivative of {{chem2|P4}} wherein one P-P bond is broken and one additional bond is formed with the neighbouring tetrahedron, resulting in chains of {{chem2|P21}} molecules linked by [[van der Waals force]]s.<ref name="Shen2016" /> Red phosphorus may be formed by heating white phosphorus to {{convert|250|C|K}} in the absence of air or by exposing it to sunlight.<ref name="Mellor1939" /> In this form phosphorus is [[amorphous]], but can be crystallised upon further heating into violet phosphorus or fibrous red phosphorus depending on the reaction conditions. Red phosphorus is therefore not an allotrope in the strictest sense of the term, but rather an intermediate between other crystalline allotropes of phosphorus, and consequently most of its properties have a range of values. Freshly prepared, bright red phosphorus is highly reactive and ignites at about {{convert|300|C|K}}.<ref name="Wiberg2001" /> After prolonged heating or storage, the color darkens; the resulting product is more stable and does not spontaneously ignite in air.<ref name="Hammond2000" />

Violet phosphorus or α-metallic phosphorus can be produced by day-long annealing of red phosphorus above {{convert|550|C|K}}. In 1865, [[Johann Wilhelm Hittorf]] discovered that when phosphorus was recrystallised from molten [[lead]], a red/purple form is obtained. Therefore, this form is sometimes known as "Hittorf's phosphorus" .<ref name="Berger1996" />

Black phosphorus or β-metallic phosphorus is the least reactive allotrope and the thermodynamically stable form below {{convert|550|C|K}}. In appearance, properties, and structure, it resembles [[graphite]], being black and flaky, a conductor of electricity, and having puckered sheets of linked atoms.<ref name="Engel2003" /><ref name="Brown1965" /><ref name="Cartz1979" /> It is obtained by heating white phosphorus under high pressures (about {{convert|12000|atm|GPa|disp=or}}). It can also be produced at ambient conditions using metal salts, e.g. mercury, as catalysts.<ref name="Lange2007" /> Single-layer black phosphorus is called [[phosphorene]], and is therefore predictably analogous to [[graphene]].

===Natural occurrence===
{{See also|Abundance of elements in Earth's crust}}
In 2013, astronomers detected phosphorus in [[Cassiopeia A]], which confirmed that this element is produced in [[supernova]]e as a byproduct of [[supernova nucleosynthesis]]. The phosphorus-to-[[iron]] ratio in material from the [[supernova remnant]] could be up to 100 times higher than in the [[Milky Way]] in general.<ref name="Koo2013" /> In 2020, astronomers analysed [[Atacama Large Millimeter Array|ALMA]] and [[Rosetta (spacecraft)#Gas and particles|ROSINA]] data from the massive [[Star formation|star-forming region]] AFGL 5142, to detect phosphorus-bearing molecules and how they could have been carried in comets to the early Earth.<ref name="Rivilla2019" />

Phosphorus has a concentration in the [[Earth's crust]] of about one gram per kilogram (for comparison, copper is found at about 0.06 grams per kilogram). It is not found free in nature, but is widely distributed in many [[mineral]]s, usually as phosphates. Inorganic [[phosphate rock]], which is partially made of [[apatite]], is today the chief commercial source of this element.

==Compounds==
{{Main category|Phosphorus compounds}}
===Inorganic phosphates===
====Phosphoric acids====
{{main|Phosphoric acids and phosphates}}
The most prevalent compounds of phosphorus are derivatives of phosphate ({{chem2|PO4(3−)}}), a tetrahedral anion.<ref name="Corbridge1995" /> Phosphate is the conjugate base of phosphoric acid, which is produced on a massive scale for use in fertilisers. Being triprotic, phosphoric acid converts stepwise to three conjugate bases:
:{{chem2|H3PO4 + H2O <-> H3O+ + H2PO4-}} (''K''<sub>a1</sub>&thinsp;= 7.25×10<sup>−3</sup>)
:{{chem2|H2PO4- + H2O <-> H3O+ + HPO4(2-)}} (''K''<sub>a2</sub>&nbsp;= 6.31×10<sup>−8</sup>)
:{{chem2|HPO4(2-) + H2O <-> H3O+ + PO4(3-)}} (''K''<sub>a3</sub>&nbsp;= 3.98×10<sup>−13</sup>)

Food-grade [[phosphoric acid]] (additive [[E number|E338]]<ref name="FGOVUK" />) is used to acidify foods and beverages such as various [[cola]]s and jams, providing a tangy or sour taste.<ref name="Threlfall1951" /> The phosphoric acid also serves as a [[preservative]].<ref name="Coca-ColaGB" /> Soft drinks containing phosphoric acid, including [[Coca-Cola]], are sometimes called [[phosphate soda]]s or phosphates. Phosphoric acid in soft drinks has the potential to cause dental erosion,<ref name="Moynihan2002" /> as well as contribute to the formation of [[Kidney stone disease|kidney stones]], especially in those who have had kidney stones previously.<ref name="Qaseem2014" />

====Metal salts====
With metal [[cation]]s, 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 ({{chem2|HPO4(2-)}}).

Calcium phosphates in particular are widespread compounds with many applications. Among them, they are used to improve the characteristics of processed meat and [[cheese]], in [[baking powder]], and in toothpaste.<ref name="Threlfall1951" /> Two of the most relevant among them are [[monocalcium phosphate]], and [[dicalcium phosphate]].

====Polyphosphates====
Phosphate exhibits a tendency to form chains and rings containing P-O-P bonds. Many polyphosphates are known, including [[Adenosine triphosphate|ATP]]. Polyphosphates arise by dehydration of hydrogen phosphates such as {{chem2|HPO4(2-)}} and {{chem2|H2PO4-}}. For example, the industrially important pentasodium triphosphate (also known as [[sodium tripolyphosphate]], STPP) is produced industrially by the megatonne by this [[condensation reaction]]:
:{{chem2|2 Na2HPO4 + NaH2PO4 -> Na5P3O10 + 2 H2O}}
Sodium triphosphate is used in laundry detergents in some countries, but banned for this use in others.<ref name="Hammond2000" /> This compound [[water softening|softens]] the water to enhance the performance of the detergents and to prevent pipe and boiler tube [[corrosion]].<ref name="Schrödter2008" />

====Oxoacids====
{{main|Phosphorus oxoacids}}
Phosphorus [[oxoacid]]s are extensive, often commercially important, and sometimes structurally complicated. They all have acidic protons bound to oxygen atoms, some have nonacidic protons that are bonded directly to phosphorus and some contain phosphorus–phosphorus bonds.<ref name="Greenwood1997" /> Although many oxoacids of phosphorus are formed, only nine are commercially important. Among them, hypophosphorous, phosphorous and orthophosphoric acid are particularly important.

{|class="wikitable"
|-
!Oxidation state!!Formula!!Name!!Acidic protons!!Compounds
|-
| +1||{{chem2|HH2PO2}}||[[hypophosphorous acid]]||1||acid, salts
|-
| +3||{{chem2|H3PO3}}||[[phosphorous acid]]<br />(phosphonic acid)||2||acid, salts
|-
| +3||{{chem2|HPO2}}||metaphosphorous acid||1||salts
|-
| +4||{{chem2|H4P2O6}}||[[hypophosphoric acid]]||4||acid, salts
|-
| +5||{{chem2|(HPO3)_{''n''}|}}||[[metaphosphoric acid]]s||''n''||salts (''n''&nbsp;=&nbsp;3,4,6)
|-
| +5||{{chem2|H(HPO3)_{''n''}OH}}||[[polyphosphoric acid]]s||''n''+2||acids, salts (''n''&nbsp;=&thinsp;1-6)
|-
| +5||{{chem2|H5P3O10}}||[[tripolyphosphoric acid]]||3||salts
|-
| +5||{{chem2|H4P2O7}}||[[pyrophosphoric acid]]||4||acid, salts
|-
| +5||{{chem2|H3PO4}}||(ortho)[[phosphoric acid]]||3||acid, salts
|}

===Other inorganic compounds===
====Oxides and sulfides====
{{main|phosphorus oxides|phosphorus sulfides}}
[[File:Phosphorus-pentoxide-3D-balls.png|thumb|right|The tetrahedral structure of {{chem2|P4O10}} and {{chem2|P4S10}}]]
[[Phosphorus pentoxide]] ({{chem2|P4O10}}) is the [[acid anhydride]] of phosphoric acid, but several intermediates between the two are known. This waxy white solid reacts vigorously with water. Similarly, [[phosphorus trioxide]] ({{chem2|P4O6}}, also called tetraphosphorus hexoxide) is the anhydride of {{chem2|P(OH)3}}, the minor tautomer of phosphorous acid. The structure of {{chem2|P4O6}} is like that of {{chem2|P4O10}} without the terminal oxide groups. Mixed oxyhalides and oxyhydrides of phosphorus(III) are almost unknown. Meanwhile, phosphorus forms a wide range of sulfides, where the phosphorus can be in P(V), P(III) or other oxidation states. However, only two of them are commercially significant. [[Phosphorus pentasulfide]] ({{chem2|P4S10}}) has a structure analogous to {{chem2|P4O10}}, and is used in the manufacture of additives and pesticides.<ref name="Heal1980" /> The three-fold symmetric [[Phosphorus sesquisulfide]] ({{chem2|P4S3}}) is used in [[strike-anywhere match]]es.

====Halides====
{{main|phosphorus halides}}
Phosphorus [[halide]]s can have as oxidation state +3 in the case of trihalides and +5 for pentahalides and [[Chalcogen#With halogens|chalcoalide]]s, but also +2 for disphosphorus tetrahalides. All four symmetrical trihalides are well known: gaseous {{chem2|PF3|link=phosphorus trifluoride}}, the yellowish liquids {{chem2|PCl3|link=phosphorus trichloride}} and {{chem2|PBr3|link=phosphorus tribromide}}, and the solid {{chem2|PI3|link=phosphorus triiodide}}. These materials are moisture sensitive, hydrolysing to give [[phosphorous acid]]. The trichloride, a common reagent used for the manufacture of pesticides, is produced by chlorination of white phosphorus. The trifluoride is produced from the trichloride by halide exchange. {{chem2|PF3}} is toxic because it binds to [[haemoglobin]].

Most phosphorus pentahalides are common compounds. {{chem2|PF5|link=phosphorus pentafluoride}} is a colourless gas and the molecules have a [[trigonal bipyramid]]al geometry. With fluoride, it forms {{chem2|PF6-}}, an [[anion]] that is [[isoelectronic]] with {{chem2|SF6|link=sulfur hexafluoride}}. {{chem2|PCl5|link=phosphorus pentachloride}} is a colourless solid which has an ionic formulation of {{chem2|PCl4+PCl6-}}, but adopts a trigonal bipyramidal geometry when molten or in the vapour phase.<ref name="Greenwood1997" /> Both the pentafluoride and the pentachloride are [[Lewis acid]]s. Meanwhile, {{chem2|PBr5|link=phosphorus pentabromide}} is an unstable solid formulated as {{chem2|PBr4+Br-}}. {{chem2|PI5|link=phosphorus pentaiodide}} is not known.<ref name="Greenwood1997" />

The most important phosphorus [[oxyhalide]] is [[phosphorus oxychloride]] ({{chem2|POCl3}}), which is approximately tetrahedral. It is prepared from {{chem2|PCl3}} and used in the manufacture of plasticizers. Phosphorus can also form thiohalides such as {{chem2|PSCl3|link=Thiophosphoryl chloride}}, and in rare cases selenohalides.

====Nitrides====
The PN molecule [[phosphorus mononitride]] is considered unstable, but is a product of crystalline [[triphosphorus pentanitride]] decomposition at {{convert|1100|K|C}}. Similarly, {{chem2|H2PN}} is considered unstable, and phosphorus nitride halogens like {{chem2|F2PN}}, {{chem2|Cl2PN}}, {{chem2|Br2PN}}, and {{chem2|I2PN}} oligomerise into cyclic [[polyphosphazene]]s. For example, compounds of the formula {{chem2|(PNCl2)_{''n''}|}} exist mainly as rings such as the [[trimer (chemistry)|trimer]] [[hexachlorophosphazene]]. The phosphazenes arise by treatment of phosphorus pentachloride with ammonium chloride:
:{{chem2|PCl5 + NH4Cl -> 1/''n'' (NPCl2)_{''n''} + 4 HCl}}
When the chloride groups are replaced by [[alkoxide]] ({{chem2|RO-}}), a family of polymers is produced with potentially useful properties.<ref name="Mark1992" />

====Phosphides and phosphine====
{{main|Phosphide|Template:Phosphides}}
A wide variety of compounds which contain the containing the phosphide ion {{chem2|P(3−)}} exist, both with [[main-group element]]s and with [[metal]]s. They often exhibit complex structures, where phosphorus has the −3 oxidation state. Metal phosphides arise by reaction of metals with red phosphorus. The [[alkali metal]]s (group 1) and [[alkaline earth metal]]s (group 2) can also form compounds such as {{chem2|Na3P7|link=sodium phosphide}}. These compounds react with water to form [[phosphine]].<ref name="Greenwood1997" /> Some phosphide minerals are also known, like {{chem2|(Fe,Ni)2P|link=Allabogdanite}} and {{chem2|(Fe,Ni)3P|link=Schreibersite}}, but they are very rare on Earth, most instances occurring in [[Iron meteorite|iron-nickel meteorite]]s.

Phosphine ({{chem2|PH3}}) and its organic derivatives are structural analogues of [[ammonia]] ({{chem2|NH3}}), but the bond angles at phosphorus are closer to 90° for phosphine and its organic derivatives. It is an ill-smelling and toxic gas, produced by hydrolysis of [[calcium phosphide]] ({{chem2|Ca3P2}}). Unlike ammonia, phosphine is oxidised by air. Phosphine is also far less basic than ammonia. Other phosphines are known which contain chains of up to nine phosphorus atoms and have the formula {{chem2|P_{''n''}H_{''n''+2}|}}.<ref name="Greenwood1997" /> The highly flammable gas [[diphosphine]] ({{chem2|P2H4}}) is an analogue of [[hydrazine]].

===Organophosphorus compounds===
{{Main|organophosphorus chemistry}}
[[File:YoshifujiR2P2.png|thumb|right|A stable diphosphene, a derivative of phosphorus(I)]]
====Phosphines, phosphites and organophosphates====
{{main|phosphaalkenes|organophosphates}}
Compounds with P-C and P-O-C bonds are often classified as organophosphorus compounds. They are widely used commercially. The {{chem2|P(3+)}} serves as a source of {{chem2|PCl3}} in routes to organophosphorus(III) compounds. For example, it is the precursor to [[triphenylphosphine]]:
:{{chem2|PCl3 + 6 Na + 3 C6H5Cl -> P(C6H5)3 + 6 NaCl}}
Treatment of phosphorus trihalides with alcohols and [[phenol]]s gives phosphites, e.g. [[triphenylphosphite]]:
:{{chem2|PCl3 + 3 C6H5OH -> P(OC6H5)3 + 3 HCl}}
Similar reactions occur for [[phosphorus oxychloride]], affording [[triphenylphosphate]]:
:{{chem2|OPCl3 + 3 C6H5OH -> OP(OC6H5)3 + 3 HCl}}

Some organophosphates are used as flame retardants.<ref name="Naiker2023" /> Among them, [[tricresyl phosphate]] and [[2-Ethylhexyl diphenyl phosphate|2-ethylhexyl diphenyl phosphate]] are also [[plasticisers]], making these two properties useful in the production of non-flammable plastic products and derivatives.<ref name="Greenwood1997" /><ref name="Diskowski2000" />

While many organic compounds of phosphorus are required for life, some are highly toxic. A wide range of organophosphorus compounds are used for their toxicity as [[pesticide]]s and [[weapon]]ised as [[nerve agent]]s.<ref name="Greenwood1997" /> Some notable examples include [[sarin]], [[VX (nerve agent)|VX]] or [[Tabun (nerve agent)|Tabun]]. Fluorophosphate [[ester]]s (like sarin) are among the most potent [[neurotoxin]]s known.

====Thioesters====
Symmetric phosphorus(III) trithioesters (e.g. {{chem2|P(SMe)3}}) can be produced from the reaction of [[white phosphorus]] and the corresponding [[disulfide]], or phosphorus(III) halides and [[thiolate]]s. Unlike the corresponding esters, they do not undergo a variant of the [[Michaelis-Arbuzov reaction]] with electrophiles.  Instead, they revert to another phosphorus(III) compound through a [[sulfonium]] intermediate.<ref name="Senning1971" />

====Phosphorus(I) and phosphorus(II)====
{{main|diphosphenes|diphosphane#organic diphosphanes}}
These compounds generally feature P–P bonds.<ref name="Greenwood1997" /> Examples include catenated derivatives of phosphine and organophosphines. Compounds containing P=P double bonds have also been observed, although they are rare.

==Biological role==
===Cells===
Inorganic phosphorus in the form of the phosphate {{chem2|PO4(3-)}} is required for all known forms of [[life]].<ref name="Ruttenberg2011" /> Phosphorus plays a major role in the structural framework of [[DNA]] and [[RNA]]. Living cells use phosphate to transport cellular energy with [[adenosine triphosphate]] (ATP), necessary for every cellular process that uses energy. ATP is also important for [[phosphorylation]], a key regulatory event in cells. Every living cell is encased in a membrane that separates it from its surroundings. Cellular membranes are composed of a phospholipid matrix and proteins, typically in the form of a bilayer. [[Phospholipid]]s are derived from [[glycerol]] with two of the glycerol hydroxyl (OH) protons replaced by fatty acids as an [[ester]], and the third hydroxyl proton has been replaced with phosphate bonded to another alcohol.<ref name="Nelson2000" />

===Bone and teeth enamel===
{{see also|Calcium metabolism}}
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 [[fluorapatite]]:<ref name="Greenwood1997" />
:{{chem2|Ca5(PO4)3OH + F- -> Ca5(PO4)3F + OH-}}
An average adult human contains about {{convert|0.7|kg|lb}} of phosphorus, about 85–90% in bones and teeth in the form of [[apatite]], and the remainder in soft tissues and extracellular fluids. The phosphorus content increases from about 0.5% by mass in infancy to 0.65–1.1% by mass in adults. In comparison, average phosphorus concentration in the blood is about 0.4&nbsp;g/L; about 70% of that is organic and 30% inorganic phosphates.<ref name="Bernhardt2008" />

===Nutrition===
The main food sources for phosphorus are the same as those containing [[protein]], although proteins themselves do not contain phosphorus. For example, milk, meat, and soya typically also have phosphorus. Generally, if a diet includes sufficient protein and calcium, the amount of phosphorus is sufficient.<ref name="Medline" />

According to the [[U.S. Institute of Medicine]], the estimated average requirement for phosphorus for people ages 19 and up is 580&nbsp;mg/day. The RDA is 700&nbsp;mg/day. RDAs are higher than EARs so as to identify amounts that will cover people with higher-than-average requirements. RDA for pregnancy and lactation are also 700&nbsp;mg/day. For people ages 1–18 years, the RDA increases with age from 460 to 1250&nbsp;mg/day. As for safety, the IOM sets [[tolerable upper intake level]] for phosphorus at 4000&nbsp;mg/day. Collectively, these values are referred to as the [[Dietary Reference Intake]].<ref name="IOM1997" /> The [[European Food Safety Authority]] (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR.<ref name="EFSA2024" /> AI and UL are defined the same as in the United States. For people ages 15 and older, including pregnancy and [[lactation]], the AI is set at 550&nbsp;mg/day. For children ages 4–10, the AI is 440&nbsp;mg/day, and for ages 11–17 it is 640&nbsp;mg/day. These AIs are lower than the U.S. RDAs. In both systems, teenagers need more than adults.<ref name="EFSA2017" /> The EFSA reviewed the same safety question and decided that there was not sufficient information to set a UL.<ref name="EFSA2006" />

Phosphorus deficiency may be caused by [[malnutrition]], by failure to absorb phosphate, and by metabolic syndromes that draw phosphate from the blood (such as in [[refeeding syndrome]] after malnutrition<ref name="Mehanna2008" />) or passing too much of it into the urine. All are characterised by [[hypophosphatemia]], which is a condition of low levels of soluble phosphate levels in the blood serum and inside the cells. Symptoms of hypophosphatemia include neurological dysfunction and disruption of muscle and blood cells due to lack of [[Adenosine triphosphate|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.<ref name="Anderson1996" />

==Phosphorus cycle==
{{Main|Phosphorus cycle}}
Phosphorus is an essential plant nutrient (the most often limiting nutrient, after nitrogen),<ref name="Etesami2019" /> and the bulk of all phosphorus production is in concentrated phosphoric acids for [[agriculture]] [[fertiliser]]s, containing as much as 70% to 75% {{chem2|P2O5}}. That led to large increase in phosphate production in the second half of the 20th century.<ref name="Philpott2013" /> Artificial phosphate fertilisation is necessary because phosphorus is essential to all living organisms; it is involved in energy transfers, strength of root and stems, [[photosynthesis]], the expansion of [[plant roots]], formation of seeds and flowers, and other important factors effecting overall plant health and genetics.<ref name="Etesami2019" /> Heavy use of phosphorus fertilisers and their runoff have resulted in [[eutrophication]] (overenrichment) of [[aquatic ecosystem]]s.<ref name="Carpenter2005" /><ref name="Conley2009" />

Natural phosphorus-bearing compounds are mostly inaccessible to plants because of the low solubility and mobility in soil.<ref name="USDA2020" /> Most phosphorus is very stable in the soil minerals or organic matter of the soil. Even when phosphorus is added in manure or fertiliser it can become fixed in the soil. Therefore, the natural [[phosphorus cycle]] is very slow. Some of the fixed phosphorus is released again over time, sustaining wild plant growth, however, more is needed to sustain intensive cultivation of crops.<ref name="PSE" /> Fertiliser is often in the form of superphosphate of lime, a mixture of calcium dihydrogen phosphate ({{chem2|Ca(H2PO4)2}}), and calcium sulfate dihydrate ({{chem2|CaSO4*2H2O}}) produced reacting sulfuric acid and water with [[calcium phosphate]].

Processing phosphate minerals with sulfuric acid for obtaining fertiliser is so important to the global economy that this is the primary industrial market for [[sulfuric acid]] and the greatest industrial use of elemental [[sulfur]].<ref name="Kogel2006" />

==Production==
===Mining===
[[File:The site of secondary mining of Phosphate rock in Nauru, 2007. Photo- Lorrie Graham (10729889683).jpg|thumb|right|Mining of phosphate rock in [[Nauru]]]]
[[File:ONCF E 1350 with phosphate train near Tamdrost.jpg|thumb|right|A phosphate train on its way to the [[port of Casablanca]] in Morocco.]]

Means of commercial phosphorus production besides mining are few because the [[phosphorus cycle]] does not include significant gas-phase transport.<ref name="Neset2011" /> The predominant source of phosphorus in modern times is phosphate rock (as opposed to the guano that preceded it).

[[Phosphate mining in the United States|US production of phosphate rock]] peaked in 1980 at 54.4 million metric tons. The United States was the world's largest producer of phosphate rock from at least 1900, up until 2006, when US production was exceeded by that of [[China]]. In 2019, the US produced 10 percent of the world's phosphate rock.<ref name="USGS2021" />

===Processing===
Most phosphorus-bearing material is for agriculture fertilisers. In this case where the standards of purity are modest, phosphorus is obtained from phosphate rock by what is called the "wet process." The minerals are treated with sulfuric acid to give [[phosphoric acid]]. Phosphoric acid is then neutralised to give various phosphate salts, which comprise fertilisers. In the wet process, phosphorus does not undergo redox.<ref name="Geeson2020" /> About five tons of [[phosphogypsum]] waste are generated per ton of phosphoric acid production. Annually, the estimated generation of phosphogypsum worldwide is 100 to 280 Mt.<ref name="Tayibi2009" />

For the use of phosphorus in drugs, detergents, and foodstuff, the standards of purity are high, which led to the development of the thermal process. In this process, phosphate minerals are converted to white phosphorus, which can be purified by distillation. The white phosphorus is then oxidised to phosphoric acid and subsequently neutralised with a base to give phosphate salts. The thermal process is conducted in a [[Submerged-arc furnace for phosphorus production|submerged-arc furnace]] which is energy intensive.<ref name="Geeson2020" /> Presently, about {{convert|1000000|ST|lk=on}} of elemental phosphorus is produced annually. [[Calcium phosphate]] (as [[Phosphorite|phosphate rock]]), mostly mined in Florida and North Africa, can be heated to 1,200–1,500&nbsp;°C with sand, which is mostly {{chem2|SiO2}}, and [[Coke (fuel)|coke]] to produce {{chem2|P4}}. The {{chem2|P4}} product, being volatile, is readily isolated:<ref name="Shriver2010" />

:{{chem2|4 Ca5(PO4)3F + 18 SiO2 + 30 C -> 3 P4 + 30 CO + 18 CaSiO3 + 2 CaF2}}
:{{chem2|2 Ca3(PO4)2 + 6 SiO2 + 10 C -> 6 CaSiO3 + 10 CO + P4}}

Side products from the thermal process include [[ferrophosphorus]], a crude form of {{chem2|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.<ref name="ERCO" />

===Reserves===
[[File:Global phosphate rock production USGS 1994-2022.png|thumb|Annual global phosphate rock production (megatonnes per yr), 1994–2022 (data from US Geological Survey)<ref name="USGS2023" />]]

Phosphorus comprises about 0.1% by mass of the [[Earth's crust]].<ref name="AGU2007" /> However, only concentrated forms collectively referred to as [[phosphate rock]] or phosphorite are exploitable, and are not evenly distributed across the Earth.<ref name="Greenwood1997" /> Unprocessed phosphate rock has a concentration of 1.7–8.7% phosphorus by mass (4–20% phosphorus pentoxide). The world's total commercial phosphate reserves and resources are estimated in amounts of phosphate rock, which in practice includes over 300 ores of different origin, composition, and phosphate content. "Reserves" refers to the amount assumed recoverable at current market prices and "resources" refers to estimated amounts of such a grade or quality that they have reasonable prospects for economic extraction.<ref name="Sutton2013" /><ref name="CIM2010" /> Mining is currently the only cost-effective method for the production of phosphorus. Hence, a shortage in rock phosphate or significant price increases might negatively affect the world's [[food security]].<ref name="Amundson2015" />

[[File:Global distribution of commercial reserves of rock phosphate USGS 2016; GTK 2015.jpg|thumb|upright 1.3|Global distribution of commercial reserves of rock phosphate in 2016]]
The countries estimated to have the biggest phosphate rock commercial reserves (in billion metric tons) are [[Morocco]] (50), [[China]] (3.2), [[Egypt]] (2.8), [[Algeria]] (2.2), [[Syria]] (1.8), [[Brazil]] (1.6), [[Saudi Arabia]] (1.4), [[South Africa]] (1.4), [[Australia]] (1.1), [[United States]] (1.0), and [[Finland]] (1.0).<ref name="Ahokas2015" /><ref name="USGS2023" /><ref name="USGS2025" /> Estimates for future production vary significantly depending on modelling and assumptions on extractable volumes, but it is inescapable that future production of phosphate rock will be heavily influenced by Morocco in the foreseeable future.<ref name="Walan2014" /> According to some researchers, Earth's commercial and affordable phosphorus reserves are expected to be depleted in 50–100 years.<ref name="Cordell2009" />

In 2023, the [[United States Geological Survey]] (USGS) estimated that economically extractable phosphate rock reserves worldwide are 72 billion tons, while world mining production in 2022 was 220 million tons.<ref name="USGS2023" /> Assuming zero growth, the reserves would thus last for around 300 years. This broadly confirms a 2010 [[International Fertilizer Development Center]] (IFDC) report that global reserves would last for several hundred years.<ref name="IDFC2010" /><ref name="VanKauwenbergh2010" /> Phosphorus reserve figures are intensely debated.<ref name="Sutton2013" /><ref name="Cordell2009" /><ref name="VanVuuren2010" /> Gilbert suggest that there has been little external verification of the estimate.<ref name="Gilbert2009" /> A 2014 review concluded that the IFDC report "presents an inflated picture of global reserves, in particular those of Morocco, where largely hypothetical and inferred resources have simply been relabeled “reserves".<ref name="Edixhoven2014" />

===Conservation and recycling===
[[File:Yorkshire Water Sewage Treatment Works (Phosphate Removal) - geograph.org.uk - 5979420.jpg|thumb|A phosphate removal sewage treatment station in [[Yorkshire]], England]]

Reducing agricultural runoff and soil erosion can slow the frequency with which farmers have to reapply phosphorus to their fields. Agricultural methods such as [[no-till farming]], [[Terrace (agriculture)|terracing]], [[contour plowing|contour tilling]], and the use of [[windbreak]]s have been shown to reduce the rate of phosphorus depletion from farmland, though do not completely remove the need for periodic fertiliser application. Strips of grassland or forest between arable land and rivers can also greatly reduce losses of phosphate and other nutrients.<ref name="Udawatta2011" />

[[Sewage treatment]] plants that have a [[Enhanced biological phosphorus removal|dedicated phosphorus removal step]] produce phosphate-rich [[sewage sludge]] that can then be [[Sewage sludge treatment|treated]] to extract phosphorus from it. This is done by [[incinerating]] the sludge and recovering the resulting ash.<ref name="Tweed2009" /> Another approach lies into the recovery of phosphorus-rich materials such as [[struvite]] from waste processing plants, which is done by adding magnesium to the waste.<ref name="Gilbert2009" /> However, the technologies currently in use are not yet cost-effective, given the current price of phosphorus on the world market.<ref name="Sartorius2011" />

==Applications==
{{About|the applications of elemental phosphorus|the applications of phosphorus compounds|Phosphorus#Compounds|section=yes}}

===Matches===
{{main|Match}}
[[File:Match striking surface.jpg|thumb|Match striking surface made of a mixture of red phosphorus, glue and ground glass. The glass powder is used to increase the friction.]]
Safety matches are very difficult to ignite on any surface other than a special striker strip. The strip contains non-toxic red phosphorus and the match head [[potassium chlorate]], an oxygen-releasing compound. When struck, small amounts of [[Abrasion (mechanical)|abrasion]] from match head and striker strip are mixed intimately to make a small quantity of [[Armstrong's mixture]], a very touch sensitive composition. The fine powder ignites immediately and provides the initial spark to set off the match head. Safety matches separate the two components of the ignition mixture until the match is struck. This is the key safety advantage as it prevents accidental ignition.<ref name="Wiberg2001" /><ref name="Hardt2001" />

===Military===
Though military uses of white phosphorus are constrained by modern international law, [[white phosphorus munition]]s are still used for military applications, such as [[incendiary device|incendiary bombs]], [[smoke screen]]s, [[smoke bomb]]s, and [[tracer ammunition]].

===Drug production===
Elemental phosphorus can reduce elemental [[iodine]] to [[hydroiodic acid]], which is a reagent effective for reducing [[ephedrine]] or [[pseudoephedrine]] to [[methamphetamine]].<ref name="Skinner1990" /> For this reason, red and white phosphorus are listed in the United States as [[DEA list of chemicals#List I chemicals|List I precursor chemicals]] by the [[Drug Enforcement Administration]], and their handling is subject to stringent regulatory controls.<ref name="CFR1" /><ref name="CFR2" /><ref name="CSA" />

===Metallurgical aspects===
Phosphorus is also an important component in [[steel]] production, in the making of [[phosphor bronze]], and in many other related products.<ref name="Scholz2014" /><ref name="Schwartz2016" /> Phosphorus is added to metallic copper during its smelting process to react with oxygen present as an impurity in copper and to produce phosphorus-containing copper ([[CuOFP]]) alloys with a higher [[hydrogen embrittlement]] resistance than normal copper.<ref name="Davisz2001" /> [[Phosphate conversion coating]] is a chemical treatment applied to steel parts to improve their corrosion resistance.

===Semiconductors===
Phosphorus is a [[doping (semiconductor)|dopant]] in [[Extrinsic semiconductor#N-type semiconductors|N-type semiconductor]]s used in high-power electronics and [[semiconductor detector]]s.<ref name="MIT" /> In this context, phosphorus is not present at the start of the process, but rather created directly out of silicon during the manufacture of the devices. This is done by neutron [[nuclear transmutation|transmutation]] doping, a method based on the conversion of the [[Isotopes_of_silicon|{{chem2|^{30}Si}}]] into {{chem2|^{31}P}} by [[neutron capture]] and [[beta decay]] as follows:
<math chem display="block">^{30}\mathrm{Si} \, (n,\gamma) \, ^{31}\mathrm{Si} \rightarrow \, ^{31}\mathrm{P} + \beta^- \; (T_{1/2} = 2.62 \mathrm{h})</math>

In practice, the silicon is typically placed near or inside a [[nuclear reactor]] generating neutrons. As neutrons pass through the silicon, phosphorus atoms are produced by transmutation. This doping method is far less common than diffusion or ion implantation, but it has the advantage of creating an extremely uniform dopant distribution.<ref name="Baliga1987" /><ref name="Schmidt1998" />

==Hazards==
===External contact===
{{Chembox
| Name = White phosphorus
| container_only = yes
| Section8 = {{Chembox Hazards
| Hazards_ref = <ref>{{Sigma-Aldrich|id=302554|name=Phosphorus, white}}</ref>
| GHSPictograms = {{GHS02}} {{GHS05}} {{GHS06}} {{GHS09}}
| GHSSignalWord = Danger
| HPhrases = {{H-phrases|H250|H300+H330|H314|H400}}
| PPhrases = {{P-phrases|P210|P222|P260|P264|P270|P271|P273|P280|P284|P301+P310+P330|P301+P330+P331|P303+P361+P353|P304+P340+P310|P305+P351+P338+P310|P335+P334|P363|P370+P378|P391|P403+P233|P405|P422|P501}}
| NFPA-H = 4
| NFPA-F = 4
| NFPA-I = 2
| TLV = 0.1 mg/m<sup>3</sup>
 }}
}}

Elemental phosphorus poses by far the greatest danger in its white form, red phosphorus being relatively nontoxic.<ref name="PublicHealthEngland" /> In the past, external exposure to white phosphorus was treated by washing the affected area with 2% [[copper(II) sulfate]] solution to form harmless compounds that are then washed away. According to 2009 [[United States Navy]] guidelines:<ref name="USNavy" />
{{blockquote|Cupric (copper) 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]].}}
Instead, the manual suggests:
{{blockquote|[...] a bicarbonate solution to neutralise 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 [[Debridement|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 [[Topical medication#Ointment|ointments]] until it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns.}}

===Ingestion===
Because of its common use as a [[rodenticide]], there are documented medical reports of white phosphorus ingestion and its effects, especially on children.<ref name="Simon1976" /> These cases can present very characteristic symptoms, such as garlic-smelling, smoking and luminescent vomit and stool, the latter sometimes called "Smoking Stool Syndrome". It is absorbed by both the gastrointestinal tract and the respiratory mucosa, to whose it causes serious damage. The acute lethal dose has been estimated at around 1&nbsp;mg/kg, this very small amount leading to many cases proving fatal, either because of rapid cardiovascular arrest or through the following systemic toxicity.<ref name="Simon1976" />

===Passive exposure===
Chronic poisoning can lead to [[phossy jaw|necrosis of the jaw]]. In the United States, exposure to 0.1&nbsp;mg/m<sup>3</sup> of white phosphorus over an 8-hour workday is set as the [[permissible exposure limit]] by the [[Occupational Safety and Health Administration]] and as the [[recommended exposure limit]] by the [[National Institute for Occupational Safety and Health]]. From 5&nbsp;mg/m<sup>3</sup>, it is considered [[immediately dangerous to life or health]].<ref name="CDC" />

==References==
{{Sister project links
 |1= |display= |auto=no |collapsible= |position=left |style=
 |wikt=phosphorus |c=yes |n=no |q=no |s=no |b=yes |v=yes
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}}

===Citations===
{{reflist|refs=
<!---In alphabetical order--->
<ref name="AGU2007">American Geophysical Union, Fall Meeting 2007, abstract #V33A-1161. [http://adsabs.harvard.edu/abs/2007AGUFM.V33A1161P Mass and Composition of the Continental Crust]</ref>

<ref name="Ahokas2015">{{cite web |last=Ahokas |first=K. |date=2015 |title=Finland's phosphorus resources are more important than ever (Geological Survey of Finland) |url=http://verkkolehti.geofoorumi.fi/en/2015/10/finlands-phosphorus-resources-are-more-important-than-ever/ |url-status=dead |archive-url=https://web.archive.org/web/20190506172330/http://verkkolehti.geofoorumi.fi/en/2015/10/finlands-phosphorus-resources-are-more-important-than-ever/ |archive-date=2019-05-06 |access-date=2017-04-01}}</ref>

<ref name="Amundson2015">{{cite journal |last1=Amundson |first1=R. |last2=Berhe |first2=A. A. |last3=Hopmans |first3=J. W. |last4=Olson |first4=C. |last5=Sztein |first5=A. E. |last6=Sparks |first6=D. L. |year=2015 |title=Soil and human security in the 21st century |url=http://www.escholarship.org/uc/item/8f42m6w4 |journal=Science |volume=348 |issue=6235 |pages=1261071 |bibcode=2015Sci...34861071A |doi=10.1126/science.1261071 |issn=0036-8075 |pmid=25954014 |s2cid=206562728}}</ref>

<ref name="Anderson1996">{{cite journal |last=Anderson |first=John J. B. |date=1996 |title=Calcium, Phosphorus and Human Bone Development |journal=[[Journal of Nutrition]] |volume=126 |issue=4 Suppl |pages=1153S–1158S |doi=10.1093/jn/126.suppl_4.1153S |pmid=8642449 |doi-access=free}}</ref>

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<!---End of reflist--->
}}

===General sources===
*{{cite book |last=March |first=Jenny |title=Oxford Research Encyclopedia of Classics |chapter=Phosphorus |publisher=Oxford University Press |date=2015-12-22 |isbn=978-0-19-938113-5 |doi=10.1093/acrefore/9780199381135.013.5052 |url=http://classics.oxfordre.com/view/10.1093/acrefore/9780199381135.001.0001/acrefore-9780199381135-e-5052 |access-date=2025-06-07 |page=}}

*{{cite book |last=Maciá-Barber |first=Enrique |title=The Chemical Evolution of Phosphorus: An Interdisciplinary Approach to Astrobiology |publisher=Apple Academic Press |publication-place=Includes bibliographical references and index. |date=2019-11-21 |isbn=978-0-429-26513-6 |doi=10.1201/9780429265136 |url=https://www.taylorfrancis.com/books/9780429555688 |access-date=2025-06-07 |page=}}

*{{cite book |last=Corbridge |first=D.E.C. |title=Phosphorus |publisher=CRC Press |date=2016-04-21 |isbn=978-1-4398-4089-4 |doi=10.1201/b12961 |url=https://www.taylorfrancis.com/books/9781439840894 |access-date=2025-06-07 |page=}}

*{{cite book |last=Francois |first=M.Rony |last2=Stephen |first2=Frank |title=Hamilton & Hardy's Industrial Toxicology |chapter=Phosphorus Compounds |publisher=John Wiley & Sons, Inc. |publication-place=Hoboken, New Jersey |date=2015-03-20 |isbn=978-1-118-83401-5 |doi=10.1002/9781118834015.ch51 |url=https://onlinelibrary.wiley.com/doi/10.1002/9781118834015.ch51 |access-date=2025-06-07 |page=383–390}}

*{{cite book |last=Walsh |first=Christopher T |title=The Chemical Biology of Phosphorus |publisher=The Royal Society of Chemistry |date=2020-10-30 |isbn=978-1-83916-202-2 |doi=10.1039/9781839162312 |url=https://books.rsc.org/books/book/916/The-Chemical-Biology-of-Phosphorus |access-date=2025-06-07 |page=}}

*{{cite book |last=Murtaugh |first=Robert J. |last2=Dhupa |first2=Nishi |last3=Ross |first3=James N. |last4=Rozanski |first4=Elizabeth |last5=Rush |first5=John |title=Critical Care |publisher=Routledge |publication-place=New York |date=2021-12-20 |isbn=978-1-315-14062-9 |doi=10.1201/9781315140629 |url=https://www.taylorfrancis.com/books/9781315140629 |access-date=2025-06-07 |page=}}

*{{cite book |title=Fundamentals and Applications of Phosphorus Nanomaterials |publisher=American Chemical Society |publication-place=Washington, DC |volume=1333 |date=2019 |isbn=978-0-8412-3658-5 |doi=10.1021/bk-2019-1333 |url=https://pubs.acs.org/doi/book/10.1021/bk-2019-1333 |access-date=2025-06-07 |page=}}


{{Periodic table (navbox)}}
{{Phosphorus compounds}}

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[[Category:Phosphorus| ]]
[[Category:Chemical elements]]
[[Category:Pnictogens]]
[[Category:Reactive nonmetals]]
[[Category:Polyatomic nonmetals]]
[[Category:Dietary minerals]]
[[Category:Pyrotechnic fuels]]