View source for Phosphorus

{{contradicts other|Abundance of elements in Earth's crust|here|Abundance |date=October 2024}}
{{About|the chemical element}}
{{Use British English|date=January 2018}}
{{Infobox phosphorus|engvar=en-GB}}

'''Phosphorus''' is a [[chemical element]]; it has [[Symbol (chemistry)|symbol]] '''P''' and [[atomic number]] 15. Elemental phosphorus exists in two major forms, [[white phosphorus]] and [[red phosphorus]], but because it is highly [[Reactivity (chemistry)|reactive]], phosphorus is never found as a free element on Earth. It has a concentration in the Earth's crust of about 0.1%, less abundant than [[hydrogen]] but more than [[manganese]]. In minerals, phosphorus generally occurs as [[phosphate]].

Elemental phosphorus was first isolated as white phosphorus in 1669. In white phosphorus, phosphorus atoms are arranged in groups of 4, written as P<sub>4</sub>. White phosphorus emits a faint glow when exposed to [[oxygen]] {{ndash}} hence, a name, taken from Greek mythology, {{lang|el|Φωσφόρος}} meaning 'light-bearer' (Latin {{lang|la|[[Lucifer]]}}), referring to the "[[Phosphorus (morning star)|Morning Star]]", the planet [[Venus]]. The term ''[[phosphorescence]]'', meaning glow after illumination, has its origin in phosphorus, although phosphorus itself does not exhibit phosphorescence: phosphorus glows due to [[oxidation]] of the white (but not red) phosphorus {{ndash}} a process now called [[chemiluminescence]]. Phosphorus is classified as a [[pnictogen]], together with [[nitrogen]], [[arsenic]], [[antimony]], [[bismuth]], and [[moscovium]]. 

Phosphorus is an element essential to sustaining [[life]] largely through [[phosphate]]s, compounds containing the phosphate ion, PO<sub>4</sub><sup>3−</sup>. Phosphates are a component of [[DNA]], [[RNA]], [[Adenosine triphosphate|ATP]], and [[phospholipid]]s, complex compounds fundamental to [[Cell (biology)|cells]]. Elemental phosphorus was first isolated from human [[urine]], and [[bone ash]] was an important early phosphate source. Phosphate mines contain fossils because phosphate is present in the fossilized deposits of animal remains and excreta. Low phosphate levels are an important limit to growth in a number of plant ecosystems. The vast majority of phosphorus compounds mined are consumed as [[fertiliser]]s. Phosphate is needed to replace the phosphorus that plants remove from the soil, and its annual demand is rising nearly twice as fast as the growth of the human population. Other applications include [[organophosphorus compound]]s in [[detergent]]s, [[pesticide]]s, and [[nerve agent]]s.

==Characteristics==
===Allotropes===
{{Main|Allotropes of phosphorus}}
Phosphorus has several [[allotropy|allotropes]] that exhibit strikingly diverse properties.<ref name="HW85">{{Cite book|author=A. Holleman|author2=N. Wiberg|title=Lehrbuch der Anorganischen Chemie|publisher= de Gruyter|date=1985|chapter=XV 2.1.3|edition= 33rd|isbn=3-11-012641-9}}</ref> The two most common allotropes are white phosphorus and red phosphorus.<ref name="Abundance">[http://www.ptable.com/#Property/Abundance/Crust Abundance]. ptable.com</ref>

For both pure and applied uses, the most important allotrope is [[Allotropes of phosphorus#White phosphorus|white phosphorus]], often abbreviated WP. White phosphorus is a soft, waxy [[molecular solid]] composed of {{chem|P|4}} [[tetrahedra]]. This {{chem|P|4}} tetrahedron is also present in liquid and gaseous phosphorus up to the temperature of {{convert|800|C|F K|-2}} when it starts decomposing to {{chem|P|2}} molecules.<ref name=WPPoly>{{Cite journal|doi=10.1002/cber.19971300911|title=On the Polymorphism of White Phosphorus|date=1997|author=Simon, Arndt|journal=Chemische Berichte|volume=130|pages=1235–1240|last2=Borrmann|first2=Horst|last3=Horakh|first3=Jörg|issue=9}}</ref> The nature of bonding in this {{chem|P|4}} 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 nA/T, much more than in the archetypical [[Aromaticity|aromatic]] molecule [[benzene]] (11 nA/T).<ref name="Cossairt-2010">{{Cite journal|last1=Cossairt|first1=Brandi M.|last2=Cummins|first2=Christopher C.|last3=Head|first3=Ashley R.|last4=Lichtenberger|first4=Dennis L.|last5=Berger|first5=Raphael J. F.|last6=Hayes|first6=Stuart A.|last7=Mitzel|first7=Norbert W.|last8=Wu|first8=Gang|date=2010-06-01|title=On the Molecular and Electronic Structures of AsP3 and P4|journal=Journal of the American Chemical Society|volume=132|issue=24|pages=8459–8465|doi=10.1021/ja102580d|pmid=20515032|bibcode=2010JAChS.132.8459C |issn=0002-7863}}</ref>

{{multiple image|perrow=2|total_width=320|caption_align=center
| header = {{font|size=100%|font=Sans-serif|text=Crystalline structures of some 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
}}

White phosphorus exists in two crystalline forms: α (alpha) and β (beta). At room temperature, the α-form is stable. It is more common, has cubic crystal structure and at {{convert|195.2|K|C}}, it transforms into β-form, which has hexagonal crystal structure. These forms differ in terms of the relative orientations of the constituent P<sub>4</sub> tetrahedra.<ref>{{cite book|title=Drinking Water Health Advisory: Munitions|author=Welford C. Roberts|author2=William R. Hartley|publisher=CRC Press, 1992|edition=illustrated|isbn=0-87371-754-6|page=399|date=1992-06-16}}</ref><ref>{{cite book|title=Topics in Phosphate Chemistry|author=Marie-Thérèse Averbuch-Pouchot|author2=A. Durif|publisher=World Scientific, 1996|isbn=981-02-2634-9|page=3|year=1996}}</ref> 

White phosphorus is the least stable, the most reactive, the most [[Volatility (chemistry)|volatile]], the least [[density|dense]] and the most toxic of the allotropes. White phosphorus gradually changes to red phosphorus, accelerated by light and heat. Samples of white phosphorus almost always contain some red phosphorus and accordingly appear yellow. For this reason, white phosphorus that is aged or otherwise impure (e.g., weapons-grade, not lab-grade WP) is also called yellow phosphorus. White phosphorus is highly [[flammable]] and [[pyrophoricity|pyrophoric]] (self-igniting) in air; it faintly glows green and blue in the dark when exposed to oxygen. The autoxidation commonly coats samples with white [[phosphorus pentoxide]] ({{chem|P|4|O|10}}): P<sub>4</sub> tetrahedra, but with oxygen inserted between the phosphorus atoms and at the vertices. White phosphorus is a [[napalm]] additive,{{citation needed|date=October 2024}} and the characteristic odour of combustion is garlicky.{{why?|reason=Which combustion product causes this?|date=October 2024}} White phosphorus is insoluble in water but soluble in carbon disulfide.<ref name="Greenwood" />

[[Thermal decomposition]] of P<sub>4</sub> at 1100&nbsp;K gives [[diphosphorus]], P<sub>2</sub>. This species is not stable as a solid or liquid. The dimeric unit contains a triple bond and is analogous to N<sub>2</sub>. It can also be generated as a transient intermediate in solution by thermolysis of organophosphorus precursor reagents.<ref>{{Cite journal|journal = [[Science (journal)|Science]]|volume = 313|issue = 5791|doi = 10.1126/science.1129630|title = Triple-Bond Reactivity of Diphosphorus Molecules|date = 2006|author = Piro, N. A.|pmid = 16946068|last2 = Figueroa|first2 = J. S.|last3 = McKellar|first3 = J. T.|last4 = Cummins|first4 = C. C.|bibcode = 2006Sci...313.1276P|pages = 1276–9 |s2cid = 27740669}}</ref> At still higher temperatures, P<sub>2</sub> dissociates into atomic P.<ref name="Greenwood" />

{| class="wikitable floatright" style="text-align:center; font-size: 95%; margin-top:1.2em; margin-left:20px"
|+ Properties of some allotropes of phosphorus<ref name="HW85" /><ref name="berger">{{Cite book|url=https://archive.org/details/semiconductormat0000berg|url-access=registration|page=[https://archive.org/details/semiconductormat0000berg/page/84 84]|title=Semiconductor materials|author=Berger, L. I.|publisher =CRC Press| date= 1996|
isbn=0-8493-8912-7}}</ref>
!Form
!white(α)
!white(β)
!red
!violet
!black
|-
!Symmetry
|Body-centred<br />cubic
|[[Triclinic]]
|[[Amorphous]]
|[[Monoclinic]]
|[[Orthorhombic]]
|-
![[Pearson symbol]]
|
|aP24
|
|mP84
|oS8
|-
![[Space group]]
|I{{overline|4}}3m
|P{{overline|1}} No.2
|
|P2/c No.13
|Cmce No.64
|-
![[Density]] (g/cm<sup>3</sup>)
|1.828
|1.88
|~2.2
|2.36
|2.69
|-
![[Band gap]] (eV)
|2.1
|
|1.8
|1.5
|0.34
|-
![[Refractive index]]
|1.8244
|
|
|2.6
|2.4
|}

[[Red phosphorus]] is polymeric in structure. It can be viewed as a derivative of P<sub>4</sub> wherein one P-P bond is broken, and one additional bond is formed with the neighbouring tetrahedron resulting in chains of P<sub>21</sub> molecules linked by [[van der Waals forces]].<ref>{{cite book|last1=Shen|first1=Z|last2=Yu|first2=JC|editor-last1=Yamashita|editor-first1=H|editor-last2=Li|editor-first2=H|title=Nanostructured Photocatalysts: Advanced Functional Materials|date=2016|chapter=Nanostructured elemental photocatalysts: Development and challenges|pages=295–312 (301)|publisher=Springer|location=Switzerland|isbn=978-3-319-26077-8}}</ref> Red phosphorus may be formed by heating white phosphorus to {{convert|250|C|F}} or by exposing white phosphorus to sunlight.<ref name="mellor-717">{{harvnb|Parkes|Mellor|1939|page=717}}</ref> Phosphorus after this treatment is [[amorphous]]. Upon further heating, this material crystallises. In this sense, red phosphorus is not an allotrope, but rather an intermediate phase between the white and violet phosphorus, and most of its properties have a range of values. For example, freshly prepared, bright red phosphorus is highly reactive and ignites at about {{convert|300|C|F}},<ref name="wiberg">{{cite book|author1=Egon Wiberg|author2=Nils Wiberg|author3=Arnold Frederick Holleman|title=Inorganic chemistry|url=https://books.google.com/books?id=Mtth5g59dEIC&pg=PA684|access-date=2011-11-19|date=2001|publisher=Academic Press|isbn=978-0-12-352651-9|pages=683–684, 689}}</ref> though it is more stable than white phosphorus, which ignites at about {{convert|30|C|F}}.<ref name="mellor-721-722">{{harvnb|Parkes|Mellor|1939|pages=721–722}}</ref> After prolonged heating or storage, the color darkens (see infobox images); the resulting product is more stable and does not spontaneously ignite in air.<ref name="CRC" />

[[Violet phosphorus]] is a form of phosphorus that can be produced by day-long annealing of red phosphorus above 550&nbsp;°C. In 1865, [[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" (or violet or α-metallic phosphorus).<ref name="berger" />

[[Black phosphorus]] is the least reactive allotrope and the thermodynamically stable form below {{convert|550|C}}. It is also known as β-metallic phosphorus and has a structure somewhat resembling that of [[graphite]].<ref name="Brown">{{Cite journal|author = A. Brown|author2 = S. Runquist|journal = Acta Crystallogr|volume = 19|date = 1965|pages = 684–685|doi = 10.1107/S0365110X65004140|title = Refinement of the crystal structure of black phosphorus|issue = 4| bibcode=1965AcCry..19..684B }}</ref><ref>{{Cite journal|author = Cartz, L.|author2 = Srinivasa, S.R.|author3 = Riedner, R.J.|author4 = Jorgensen, J.D.|author5 = Worlton, T.G.|journal = Journal of Chemical Physics|date = 1979|volume = 71|pages = 1718–1721|doi = 10.1063/1.438523|title = Effect of pressure on bonding in black phosphorus|bibcode = 1979JChPh..71.1718C|issue = 4 }}</ref> 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>{{Cite journal|author = Lange, Stefan|author2 = Schmidt, Peer|author3 = Nilges, Tom|name-list-style = amp |journal = [[Inorg. Chem.]]|date = 2007|volume = 46|issue = 10|pmid = 17439206|doi = 10.1021/ic062192q|title = Au3SnP7@Black Phosphorus: An Easy Access to Black Phosphorus|pages = 4028–35}}</ref> In appearance, properties, and structure, it resembles [[graphite]], being black and flaky, a conductor of electricity, and has puckered sheets of linked atoms.<ref>{{cite book|title= Synthesis of Carbon-Phosphorus Bonds|author=Robert Engel|publisher=CRC Press, 2003|edition=2|isbn=0-203-99824-3|page=11|date=2003-12-18}}</ref>

Another form, scarlet phosphorus, is obtained by allowing a solution of white phosphorus in [[carbon disulfide]] to evaporate in [[sunlight]].<ref name="berger" />

===Chemiluminescence===
[[File:White phosphorus glowing e17.png|right|upright=0.9|thumb|White phosphorus exposed to air glows in the dark.]]
When first isolated, it was observed that the green glow emanating from white phosphorus 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, this process is caused by the phosphorus reacting with oxygen in the air; in a sealed container, this process will eventually stop when all the oxygen in the container is consumed. By the 18th century, it was known that in pure oxygen, phosphorus does not glow at all;<ref>{{cite web|url = https://www.nobelprize.org/prizes/chemistry/1956/ceremony-speech/|title = Nobel Prize in Chemistry 1956 – Presentation Speech by Professor A. Ölander (committee member)| access-date = 2009-05-05}}</ref> there is only a range of [[partial pressure]]s at which it does. Heat can be applied to drive the reaction at higher pressures.<ref>{{cite web| url =http://www.lateralscience.co.uk/phos/index.html| title =Phosphorus |website=Lateral Science| access-date =2009-05-05| archive-url =https://web.archive.org/web/20090221031316/http://www.lateralscience.co.uk/phos/index.html| archive-date =2009-02-21|url-status=live}}</ref>

In 1974, the glow was explained by R. J. van Zee and A. U. Khan.<ref name="shockinghistory">{{harvnb|Emsley|2000}}</ref><ref>{{cite journal|doi=10.1021/j100561a021|title=The phosphorescence of phosphorus|journal=The Journal of Physical Chemistry|volume=80|issue=20|pages=2240–2242|year=1976|last1=Vanzee|first1=Richard J.|last2=Khan|first2=Ahsan U.}}</ref> A reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived molecules HPO and {{chem|P|2|O|2}} that both emit visible light. The reaction is slow and only very little of the intermediates are required to produce the luminescence, hence the extended time the glow continues in a stoppered jar.

Since its discovery, ''[[phosphor]]s'' and ''[[phosphorescence]]'' were used loosely to describe substances that shine in the dark without burning. Although the term [[phosphorescence]] is derived from phosphorus, the reaction that gives phosphorus its glow is properly called [[chemiluminescence]] (glowing due to a cold chemical reaction), not phosphorescence (re-emitting light that previously fell onto a substance and excited it).<ref name="sommers">{{harvnb|Sommers|2007}}, [https://archive.org/details/phosphorus0000somm/page/25 p.&nbsp;25]</ref>

===Isotopes===
{{Main|Isotopes of phosphorus}}
There are 22 known [[isotopes]] of phosphorus,{{NUBASE2016|ref}} ranging from {{chem|26|P}} to {{chem|47|P}}.<ref name="neufcourt">{{cite journal |last1=Neufcourt |first1=L. |last2=Cao |first2=Y. |last3=Nazarewicz |first3=W. |last4=Olsen |first4=E. |last5=Viens |first5=F. |title=Neutron drip line in the Ca region from Bayesian model averaging |date=2019 |journal=Physical Review Letters |volume=122 |issue=6 |pages=062502–1–062502–6 |doi=10.1103/PhysRevLett.122.062502 |pmid=30822058 |arxiv=1901.07632|bibcode=2019PhRvL.122f2502N |s2cid=73508148 }}</ref> Only {{chem|31|P}} is stable and is therefore present at 100% abundance. The half-integer [[nuclear spin]] and high abundance of <sup>31</sup>P make [[phosphorus-31 NMR]] 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. These are:
* {{chem|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. 
* {{chem|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 {{chem|32|P}} penetrate skin and [[cornea]]s and any {{chem|32|P}} ingested, inhaled, or absorbed is readily incorporated into bone and [[nucleic acid]]s. For these reasons, [[Occupational Safety and Health Administration]] in the United States, and similar institutions in other developed countries require personnel working with {{chem|32|P}} to wear lab coats, disposable gloves, and safety glasses or goggles to protect the eyes, and avoid working directly over open containers. [[Biomonitoring|Monitoring]] personal, clothing, and surface contamination is also required. [[Radiation protection|Shielding]] requires special consideration. 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 shielded with low density materials such as acrylic or other plastic, water, or (when transparency is not required), even wood.<ref>{{cite web |title=Phosphorus-32 |url=http://www.oseh.umich.edu/pdf/TrainP32.pdf |publisher=University of Michigan Department of Occupational Safety & Environmental Health |access-date=2010-11-18 |archive-url=https://web.archive.org/web/20160528091951/http://www.oseh.umich.edu/pdf/TrainP32.pdf |archive-date=2016-05-28 |url-status=dead }}</ref>

==Occurrence==
{{Category see also|Phosphate minerals}}

===Universe===
In 2013, astronomers detected phosphorus in [[Cassiopeia A|Cassiopeia&nbsp;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>{{Cite journal | last1 = Koo | first1 = B.-C. | last2 = Lee | first2 = Y.-H. | last3 = Moon | first3 = D.-S. | last4 = Yoon | first4 = S.-C. | last5 = Raymond | first5 = J. C. | title = Phosphorus in the Young Supernova Remnant Cassiopeia A | doi = 10.1126/science.1243823 | journal = Science | volume = 342 | issue = 6164 | pages = 1346–8 | year = 2013 | pmid = 24337291|arxiv = 1312.3807 |bibcode = 2013Sci...342.1346K | s2cid = 35593706 }}</ref>

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 are carried in comets to the early Earth.<ref name="ARX-2019">{{cite journal | last1 = Rivilla | first1 = V. M. | last2 = Drozdovskaya | first2 = M. N. | last3 = Altwegg | first3 = K. |author3-link=Kathrin Altwegg| last4 = Caselli | first4 = P.|author4-link=Paola Caselli | last5 = Beltrán | first5 = M. T. | last6 = Fontani | first6 =  F. | last7 = van der Tak | first7 = F. F. S. | last8 = Cesaroni | first8 = R. | last9 = Vasyunin | first9 = A. | last10 = Rubin | first10 = M. | last11 = Lique | first11 = F. | last12 = Marinakis | first12 = S. | last13 = Testi | first13 = L. |title=ALMA and ROSINA detections of phosphorus-bearing molecules: the interstellar thread between star-forming regions and comets| journal = Monthly Notices of the Royal Astronomical Society | volume = 492 | pages = 1180–1198 |date=2019 |arxiv=1911.11647 | doi = 10.1093/mnras/stz3336 | doi-access = free | s2cid = 208290964 }}</ref><ref name="PHYS-20200115">{{cite news |author=ESO |title=Astronomers reveal interstellar thread of one of life's building blocks |url=https://phys.org/news/2020-01-astronomers-reveal-interstellar-thread-life.html |date=15 January 2020 |work=[[Phys.org]] |access-date=16 January 2020 }}</ref>

===Crust and organic sources===
Phosphorus has a concentration in the Earth's crust of about one gram per kilogram (compare copper at about 0.06 grams). It is not found free in nature, but is widely distributed in many [[mineral]]s, usually as phosphates.<ref name="Abundance" /> Inorganic [[phosphate rock]], which is partially made of [[apatite]] (a group of minerals being, generally, pentacalcium triorthophosphate fluoride (hydroxide)), is today the chief commercial source of this element. According to the [[US Geological Survey|US Geological Survey (USGS)]], about 50 percent of the global phosphorus reserves are in [[Berber people|Amazigh]] nations like [[Morocco]], [[Algeria]] and [[Tunisia]].<ref name="USGS">{{cite web| access-date = 2009-06-06| publisher = USGS| url = http://minerals.usgs.gov/minerals/pubs/commodity/phosphate_rock/|title = Phosphate Rock: Statistics and Information}}</ref> 85% of Earth's known reserves are in [[Morocco]] with smaller deposits in [[People's Republic of China|China]], [[Russia]],<ref name="MJ" /> [[Florida]], [[Idaho]], [[Tennessee]], [[Utah]], and elsewhere.<ref>Klein, Cornelis and Cornelius S. Hurlbut, Jr., ''Manual of Mineralogy'', Wiley, 1985, 20th ed., p. 360, {{ISBN|0-471-80580-7}}</ref> [[Albright and Wilson]] in the UK and their [[Niagara Falls]] plant, for instance, were using phosphate rock in the 1890s and 1900s from Tennessee, Florida, and the [[Îles du Connétable]] ([[guano]] island sources of phosphate); by 1950, they were using phosphate rock mainly from Tennessee and North Africa.<ref name="threlfall-51">{{harvnb|Threlfall|1951|page=51}}</ref>

Organic sources, namely [[urine]], [[bone ash]] and (in the latter 19th century) [[guano]], were historically of importance but had only limited commercial success.<ref>{{harvnb|Toy|1975}}, [https://archive.org/details/chemistryofphosp0003toya/page/389 p.&nbsp;389]</ref> As urine contains phosphorus, it has fertilising qualities which are still harnessed today in some countries, including [[Sweden]], using methods for [[reuse of excreta]]. To this end, urine can be used as a fertiliser in its pure form or part of being mixed with water in the form of [[sewage]] or [[sewage sludge]].

==Compounds==
{{See also|Phosphorus halides|Category:Phosphate minerals|Category:Phosphorus compounds}}

===Phosphorus(V)===
[[File:Phosphorus-pentoxide-3D-balls.png|thumb|right|The tetrahedral structure of P<sub>4</sub>O<sub>10</sub> and P<sub>4</sub>S<sub>10</sub>]]
The most prevalent compounds of phosphorus are derivatives of phosphate (PO<sub>4</sub><sup>3−</sup>), a tetrahedral anion.{{sfn|Corbridge|1995}} 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:
:H<sub>3</sub>PO<sub>4</sub> + H<sub>2</sub>O {{eqm}} H<sub>3</sub>O<sup>+</sup> + H<sub>2</sub>PO<sub>4</sub><sup>−</sup> &nbsp;&nbsp;&nbsp;&nbsp;&nbsp; ''K''<sub>a1</sub>&thinsp;= 7.25×10<sup>−3</sup>

:H<sub>2</sub>PO<sub>4</sub><sup>−</sup> + H<sub>2</sub>O {{eqm}} H<sub>3</sub>O<sup>+</sup> + HPO<sub>4</sub><sup>2−</sup> &nbsp;&nbsp;&nbsp;&nbsp;&nbsp; ''K''<sub>a2</sub>&nbsp;= 6.31×10<sup>−8</sup>

:HPO<sub>4</sub><sup>2−</sup> + H<sub>2</sub>O {{eqm}} H<sub>3</sub>O<sup>+</sup> + &nbsp;PO<sub>4</sub><sup>3−</sup> &nbsp;&nbsp;&nbsp;&nbsp;&nbsp;&nbsp; ''K''<sub>a3</sub>&nbsp;= 3.98×10<sup>−13</sup>

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 HPO<sub>4</sub><sup>2−</sup> and H<sub>2</sub>PO<sub>4</sub><sup>−</sup>. For example, the industrially important pentasodium triphosphate (also known as [[sodium tripolyphosphate]], STPP) is produced industrially by the megatonne by this [[condensation reaction]]:
:2 Na<sub>2</sub>HPO<sub>4</sub>  +  NaH<sub>2</sub>PO<sub>4</sub>   →   Na<sub>5</sub>P<sub>3</sub>O<sub>10</sub>  +  2 H<sub>2</sub>O
[[Phosphorus pentoxide]] (P<sub>4</sub>O<sub>10</sub>) is the [[acid anhydride]] of phosphoric acid, but several intermediates between the two are known. This waxy white solid reacts vigorously with water.

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 (HPO<sub>4</sub><sup>2−</sup>).

[[phosphorus pentachloride|PCl<sub>5</sub>]] and [[phosphorus pentafluoride|PF<sub>5</sub>]] are common compounds. PF<sub>5</sub> is a colourless gas and the molecules have [[trigonal bipyramid]]al geometry. PCl<sub>5</sub> is a colourless solid which has an ionic formulation of PCl<sub>4</sub><sup>+</sup> PCl<sub>6</sub><sup>−</sup>, but adopts the [[trigonal bipyramid]]al geometry when molten or in the vapour phase.<ref name="Greenwood" /> [[phosphorus pentabromide|PBr<sub>5</sub>]] is an unstable solid formulated as PBr<sub>4</sub><sup>+</sup>Br<sup>−</sup>and [[phosphorus pentaiodide|PI<sub>5</sub>]] is not known.<ref name="Greenwood" /> The pentachloride and pentafluoride are [[Lewis acid]]s. With fluoride, PF<sub>5</sub> forms PF<sub>6</sub><sup>−</sup>, an [[anion]] that is [[isoelectronic]] with SF<sub>6</sub>. The most important oxyhalide is [[phosphorus oxychloride]], (POCl<sub>3</sub>), which is approximately tetrahedral.

Before extensive computer calculations were feasible, it was thought that bonding in phosphorus(V) compounds involved ''d'' orbitals. Computer modeling of [[molecular orbital theory]] indicates that this bonding involves only s- and p-orbitals.<ref>{{Cite journal|author = Kutzelnigg, W.|title = Chemical Bonding in Higher Main Group Elements|url = http://web.uvic.ca/~chem421/ACIE_1984_Kutzelnigg_review.pdf|doi = 10.1002/anie.198402721|journal = Angew. Chem. Int. Ed. Engl.|volume = 23|pages = 272–295|date = 1984|issue = 4|access-date = 2009-05-24|archive-date = 2020-04-16|archive-url = https://web.archive.org/web/20200416103206/http://web.uvic.ca/~chem421/ACIE_1984_Kutzelnigg_review.pdf}}</ref>

===Phosphorus(III)===
All four symmetrical trihalides are well known: gaseous [[phosphorus trifluoride|PF<sub>3</sub>]], the yellowish liquids [[phosphorus trichloride|PCl<sub>3</sub>]] and [[phosphorus tribromide|PBr<sub>3</sub>]], and the solid [[phosphorus triiodide|PI<sub>3</sub>]]. These materials are moisture sensitive, hydrolysing to give [[phosphorous acid]]. The trichloride, a common reagent, is produced by chlorination of white phosphorus:
:P<sub>4</sub> + 6 Cl<sub>2</sub> → 4 PCl<sub>3</sub>
The trifluoride is produced from the trichloride by halide exchange. PF<sub>3</sub> is toxic because it binds to [[haemoglobin]].

[[Phosphorus trioxide|Phosphorus(III) oxide]], P<sub>4</sub>O<sub>6</sub> (also called tetraphosphorus hexoxide) is the anhydride of P(OH)<sub>3</sub>, the minor tautomer of phosphorous acid. The structure of P<sub>4</sub>O<sub>6</sub> is like that of P<sub>4</sub>O<sub>10</sub> without the terminal oxide groups.

Symmetric phosphorus(III) trithioesters (e.g. P(SMe)<sub>3</sub>) 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 reverting to another phosphorus(III) compound through a [[sulfonium]] intermediate.<ref>{{cite book|title=Sulfur in Organic and Inorganic Chemistry|volume=1|editor-first=Alexander|editor-last=Senning|year=1971|publisher=Marcel Dekker|location=New York|lccn=70-154612|isbn=0-8247-1615-9|first=Lucreţia|last=Almasi|chapter=The Sulfur&ndash;Phosphorus Bond|pages=49–55}}</ref>

===Phosphorus(I) and phosphorus(II)===
[[File:YoshifujiR2P2.png|thumb|right|A stable [[diphosphene]], a derivative of phosphorus(I)]]
These compounds generally feature P–P bonds.<ref name="Greenwood">Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. {{ISBN|0-7506-3365-4}}.</ref> Examples include catenated derivatives of phosphine and organophosphines. Compounds containing P=P double bonds have also been observed, although they are rare.

===Phosphides and phosphines===
[[Phosphide]]s arise by reaction of metals with red phosphorus. The alkali metals (group 1) and alkaline earth metals can form ionic compounds containing the [[phosphide]] ion, P<sup>3−</sup>. These compounds react with water to form [[phosphine]]. Other [[phosphide]]s, for example Na<sub>3</sub>P<sub>7</sub>, are known for these reactive metals. With the transition metals as well as the monophosphides there are metal-rich phosphides, which are generally hard refractory compounds with a metallic lustre, and phosphorus-rich phosphides which are less stable and include semiconductors.<ref name="Greenwood" /> [[Schreibersite]] is a naturally occurring metal-rich phosphide found in meteorites. The structures of the metal-rich and phosphorus-rich phosphides can be complex.

[[Phosphine]] (PH<sub>3</sub>) and its organic derivatives (PR<sub>3</sub>) are structural analogues of ammonia (NH<sub>3</sub>), but the bond angles at phosphorus are closer to 90° for phosphine and its organic derivatives. Phosphine is an ill-smelling, toxic gas. Phosphorus has an oxidation number of −3 in phosphine. Phosphine is produced by hydrolysis of [[calcium phosphide]], Ca<sub>3</sub>P<sub>2</sub>. 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 P<sub>''n''</sub>H<sub>''n''+2</sub>.<ref name="Greenwood" /> The highly flammable gas [[diphosphine]] (P<sub>2</sub>H<sub>4</sub>) is an analogue of [[hydrazine]].

===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="Greenwood" /> Although many oxoacids of phosphorus are formed, only nine are commercially important, and three of them, [[hypophosphorous acid]], [[phosphorous acid]], and [[phosphoric acid]], are particularly important.

{|class="wikitable"
|-
!Oxidation state!!Formula!!Name!!Acidic protons!!Compounds
|-
| +1||HH<sub>2</sub>PO<sub>2</sub>||[[hypophosphorous acid]] || 1 ||acid, salts
|-
| +3||H<sub>3</sub>PO<sub>3</sub>||[[phosphorous acid]]<br />(phosphonic acid) || 2 ||acid, salts
|-
| +3||HPO<sub>2</sub>||metaphosphorous acid || 1 ||salts
|-
| +4||H<sub>4</sub>P<sub>2</sub>O<sub>6</sub>||[[hypophosphoric acid]]||4||acid, salts
|-
| +5||(HPO<sub>3</sub>)<sub>''n''</sub>||[[metaphosphoric acid]]s || ''n'' ||salts (''n''&nbsp;=&nbsp;3,4,6)
|-
| +5||H(HPO<sub>3</sub>)<sub>''n''</sub>OH||[[polyphosphoric acid]]s || ''n''+2 ||acids, salts (''n''&nbsp;=&thinsp;1-6)
|-
| +5||H<sub>5</sub>P<sub>3</sub>O<sub>10</sub>||[[tripolyphosphoric acid]] || 3 || salts
|-
| +5||H<sub>4</sub>P<sub>2</sub>O<sub>7</sub>||[[pyrophosphoric acid]] || 4 ||acid, salts
|-
| +5||H<sub>3</sub>PO<sub>4</sub>||(ortho)[[phosphoric acid]] || 3 ||acid, salts
|}

===Nitrides===
The PN molecule is considered unstable, but is a product of crystalline [[triphosphorus pentanitride|phosphorus nitride]] decomposition at 1100&nbsp;K. Similarly, H<sub>2</sub>PN is considered unstable, and phosphorus nitride halogens like F<sub>2</sub>PN, Cl<sub>2</sub>PN, Br<sub>2</sub>PN, and I<sub>2</sub>PN oligomerise into cyclic [[polyphosphazene]]s. For example, compounds of the formula (PNCl<sub>2</sub>)<sub>''n''</sub> exist mainly as rings such as the [[trimer (chemistry)|trimer]] [[hexachlorophosphazene]]. The phosphazenes arise by treatment of phosphorus pentachloride with ammonium chloride:<blockquote>PCl<sub>5</sub> + NH<sub>4</sub>Cl → 1/''n'' (NPCl<sub>2</sub>)<sub>''n''</sub> + 4 HCl</blockquote>When the chloride groups are replaced by [[alkoxide]] (RO<sup>−</sup>), a family of polymers is produced with potentially useful properties.<ref>Mark, J. E.; Allcock, H. R.; West, R. "Inorganic Polymers" Prentice Hall, Englewood, NJ: 1992. {{ISBN|0-13-465881-7}}.</ref>

===Sulfides===
{{main|phosphorus sulfide}}
Phosphorus forms a wide range of sulfides, where the phosphorus can be in P(V), P(III) or other oxidation states. The three-fold symmetric P<sub>4</sub>S<sub>3</sub> is used in strike-anywhere matches. P<sub>4</sub>S<sub>10</sub> and P<sub>4</sub>O<sub>10</sub> have analogous structures.<ref>Heal, H. G. "The Inorganic Heterocyclic Chemistry of Sulfur, Nitrogen, and Phosphorus" Academic Press: London; 1980. {{ISBN|0-12-335680-6}}.</ref> Mixed oxyhalides and oxyhydrides of phosphorus(III) are almost unknown.

===Organophosphorus compounds===
{{Main|organophosphorus compounds}}
Compounds with P-C and P-O-C bonds are often classified as organophosphorus compounds. They are widely used commercially. The PCl<sub>3</sub> serves as a source of P<sup>3+</sup> in routes to organophosphorus(III) compounds. For example, it is the precursor to [[triphenylphosphine]]:
:PCl<sub>3</sub> + 6 Na + 3 C<sub>6</sub>H<sub>5</sub>Cl → P(C<sub>6</sub>H<sub>5</sub>)<sub>3</sub> + 6 NaCl
Treatment of phosphorus trihalides with alcohols and [[phenol]]s gives phosphites, e.g. [[triphenylphosphite]]:
:PCl<sub>3</sub> + 3 C<sub>6</sub>H<sub>5</sub>OH → P(OC<sub>6</sub>H<sub>5</sub>)<sub>3</sub> + 3 HCl
Similar reactions occur for [[phosphorus oxychloride]], affording [[triphenylphosphate]]:
:OPCl<sub>3</sub> + 3 C<sub>6</sub>H<sub>5</sub>OH → OP(OC<sub>6</sub>H<sub>5</sub>)<sub>3</sub> + 3 HCl

==History ==
===Etymology===
The name ''Phosphorus'' in Ancient Greece was the name for the planet [[Venus]] and is derived from the [[Greek language|Greek]] words (φῶς = light, φέρω = carry), which roughly translates as light-bringer or light carrier.<ref name="mellor-717" /> (In [[Greek mythology]] and tradition, Augerinus (Αυγερινός = morning star, still in use today), Hesperus or Hesperinus (΄Εσπερος or Εσπερινός or Αποσπερίτης = evening star, still in use today) and Eosphorus (Εωσφόρος = dawnbearer, not in use for the planet after Christianity) are close homologues, and also associated with [[Phosphorus (morning star)|Phosphorus-the-morning-star]]).

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> valence: so, just as [[sulfur]] forms ''sulfurous'' and ''sulfuric'' compounds, phosphorus forms phosphorous compounds (e.g., [[phosphorous acid]]) and P<sup>5+</sup> valence phosphoric compounds (e.g., [[phosphoric acids and phosphates]]).

===Discovery===
[[File:Robert boyle.jpg|thumb|upright|[[Robert Boyle]]]]
The discovery of phosphorus, the first element to be discovered that was not known since ancient times,<ref>{{cite journal | doi = 10.1021/ed009p11| title = The discovery of the elements. II. Elements known to the alchemists| journal = Journal of Chemical Education| volume = 9| issue = 1| page = 11| date = 1932| last1 = Weeks| first1 = Mary Elvira| bibcode = 1932JChEd...9...11W}}</ref> is credited to the German alchemist [[Hennig Brand]] in 1669, although others might have discovered phosphorus around the same time.<ref name="disc">{{harvnb|Beatty|2000}}, [https://books.google.com/books?id=FHJIUJM1_JUC&pg=PA7 p.&nbsp;7]</ref> Brand experimented with [[urine]], which contains considerable quantities of dissolved phosphates from normal metabolism.<ref name="mellor-717" /> Working in [[Hamburg]], Brand attempted to create the fabled [[philosopher's stone]] through the [[distillation]] of some [[salt (chemistry)|salt]]s by evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. It was named ''phosphorus mirabilis'' ("miraculous bearer of light").<ref name="schmundt">Schmundt, Hilmar (21 April 2010), [http://www.spiegel.de/international/world/0,1518,690450-2,00.html "Experts Warn of Impending Phosphorus Crisis"], ''[[Der Spiegel]]''.</ref>

Brand's process originally involved letting urine stand for days until it gave off a terrible stench. Then he boiled it down to a paste, heated this paste to a high temperature, and led the vapours through water, where he hoped they would condense to gold. Instead, he obtained a white, waxy substance that glowed in the dark. Brand had discovered phosphorus. Specifically, Brand produced ammonium sodium hydrogen phosphate, {{chem|(NH|4|)NaHPO|4}}. While the quantities were essentially correct (it took about {{convert|1,100|L|gal|disp=x| [|]}} of urine to make about 60&nbsp;g of phosphorus), it was unnecessary to allow the urine to rot first. Later scientists discovered that fresh urine yielded the same amount of phosphorus.<ref name="sommers" />

Brand at first tried to keep the method secret,<ref>{{Cite book| first=J. M. |last=Stillman|title = The Story of Alchemy and Early Chemistry |location = New York|publisher = Dover|date = 1960|pages = 418–419| isbn = 0-7661-3230-7}}</ref> but later sold the recipe for 200 thalers to Johann Daniel Kraft <small>([[:de:Johann Daniel Kraft|de]])</small> from Dresden.<ref name="mellor-717" /> Kraft toured much of Europe with it, including England, where he met with [[Robert Boyle]]. The secret—that the substance was made from urine—leaked out, and [[Johann von Löwenstern-Kunckel|Johann Kunckel]] (1630–1703) was able to reproduce it in Sweden (1678). Later, Boyle in London (1680) also managed to make phosphorus, possibly with the aid of his assistant, [[Ambrose Godfrey|Ambrose Godfrey-Hanckwitz]]. Godfrey later made a business of the manufacture of phosphorus.

Boyle states that Kraft gave him no information as to the preparation of phosphorus other than that it was derived from "somewhat that belonged to the body of man". This gave Boyle a valuable clue, so that he, too, managed to make phosphorus, and published the method of its manufacture.<ref name="mellor-717" /> Later he improved Brand's process by using sand in the reaction (still using urine as base material),

: 4 {{chem|NaPO|3}} + 2 {{chem|SiO|2}} + 10 C → 2 {{chem|Na|2|SiO|3}} + 10 CO + {{chem|P|4}}

Robert Boyle was the first to use phosphorus to ignite sulfur-tipped wooden splints, forerunners of modern matches, in 1680.<ref>{{cite book|title=Metabolism of the Anthroposphere|first =Peter|last= Baccini|author2=Paul H. Brunner|publisher=MIT Press, 2012|isbn=978-0-262-30054-4|page=288|date =2012-02-10}}</ref>

Phosphorus was the 13th element to be discovered. Because of its tendency to spontaneously combust when left alone in air, it is sometimes referred to as "the Devil's element".{{sfn|Emsley|2002}}

===Bone ash and guano===
[[File:DSCN5766-guano-glantz crop b.jpg|thumb|upright=0.7|[[Guano]] mining in the Central [[Chincha Islands]], c. 1860]]
[[Antoine Lavoisier]] recognized phosphorus as an element in 1777 after [[Johan Gottlieb Gahn]] and [[Carl Wilhelm Scheele]], in 1769, showed that [[calcium phosphate]] ({{chem|Ca|3|(PO|4|)|2}}) is found in bones by obtaining elemental phosphorus from [[bone ash]].<ref name=LavMem1rySrc />

Bone ash was the major source of phosphorus until the 1840s. The method started by roasting bones, then employed the use of [[fire clay]] [[retorts]] encased in a very hot brick furnace to distill out the highly toxic elemental phosphorus product.<ref>{{cite book|author=Thomson, Robert Dundas |title=Dictionary of chemistry with its applications to mineralogy, physiology and the arts|url=https://books.google.com/books?id=1LxBAAAAcAAJ&pg=PA416|year=1870|publisher=Rich. Griffin and Company|page=416}}</ref> Alternately, precipitated phosphates could be made from ground-up bones that had been de-greased and treated with strong acids. White phosphorus could then be made by heating the precipitated phosphates, mixed with ground coal or [[charcoal]] in an iron pot, and distilling off phosphorus vapour in a [[retort]].<ref name="threlfall-v">{{harvnb|Threlfall|1951|pages=49–66}}</ref> [[Carbon monoxide]] and other flammable gases produced during the reduction process were burnt off in a [[Gas flare|flare stack]].

In the 1840s, world phosphate production turned to the mining of tropical island deposits formed from bird and bat [[guano]] (see also [[Guano Islands Act]]). These became an important source of phosphates for fertiliser in the latter half of the 19th century.{{citation needed|date=October 2024}}

===Phosphate rock===
[[Phosphate rock]], which usually contains calcium phosphate, was first used in 1850 to make phosphorus, and following the introduction of the electric arc furnace by [[James Burgess Readman]] in 1888{{sfn|Toy|1975}} (patented 1889),<ref>US patent 417943</ref> elemental phosphorus production switched from the bone-ash heating, to electric arc production from phosphate rock. 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. Phosphate rock remains a feedstock in the fertiliser industry, where it is treated with sulfuric acid to produce various "[[superphosphate]]" fertiliser products.

===Incendiaries===
White phosphorus was first made commercially in the 19th century for the [[match]] industry. This used bone ash for a phosphate source, as described above. The bone-ash process became obsolete when the [[submerged-arc furnace for phosphorus production]] was introduced to reduce phosphate rock.<ref name="threlfall-vii">{{harvnb|Threlfall|1951|pages=81–101}}</ref><ref name="mellor-718-720">{{harvnb|Parkes|Mellor|1939|page=718–720}}.</ref> 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-xi">{{harvnb|Threlfall|1951|pages=167–185}}</ref> In [[World War I|World War&nbsp;I]], it was used in incendiaries, [[smoke screen]]s and tracer bullets.<ref name="threlfall-xi" /> A special incendiary bullet was developed to shoot at [[hydrogen]]-filled [[Zeppelin]]s over Britain (hydrogen being highly [[flammable]]).<ref name="threlfall-xi" /> During [[World War II|World War&nbsp;II]], [[Molotov cocktail]]s made of phosphorus dissolved in [[petrol]] 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.<ref name="Greenwood" />

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's giving off luminous steam).<ref name="shockinghistory" /> In addition, exposure to the vapours gave match workers a severe [[necrosis]] of the bones of the jaw, known as "[[phossy jaw]]". When a safe process for manufacturing red phosphorus was discovered, with its far lower flammability and toxicity, laws were enacted, under the [[Berne Convention (1906)]], requiring its adoption as a safer alternative for match manufacture.<ref name="phbook">{{Cite book| pages= 1486–1489| url =https://books.google.com/books?id=cvJuLqBxGUcC&pg=PA1487| title = Goldfrank's toxicologic emergencies| author=Lewis R. Goldfrank| author2=Neal Flomenbaum| author3=Mary Ann Howland| author4=Robert S. Hoffman| author5=Neal A. Lewin| author6=Lewis S. Nelson| publisher = McGraw-Hill Professional| date = 2006| isbn = 0-07-143763-0}}</ref> The toxicity of white phosphorus led to discontinuation of its use in matches.<ref>The White Phosphorus Matches Prohibition Act, 1908.</ref> The Allies used phosphorus [[incendiary bomb]]s in [[World War II]] to destroy Hamburg, the place where the "miraculous bearer of light" was first discovered.<ref name="schmundt" />

==Production==
[[File:The site of secondary mining of Phosphate rock in Nauru, 2007. Photo- Lorrie Graham (10729889683).jpg|thumb|upright=0.9|Mining of phosphate rock in [[Nauru]]]]

In 2017, the USGS estimated 68&nbsp;billion tons of world reserves, where reserve figures refer to the amount assumed recoverable at current market prices; 0.261&nbsp;billion tons were mined in 2016.<ref>{{cite web|url=https://minerals.usgs.gov/minerals/pubs/commodity/phosphate_rock/mcs-2017-phosp.pdf|title=Phosphate Rock|publisher=USGS|access-date=2017-03-20}}</ref> Critical to contemporary agriculture, its annual demand is rising nearly twice as fast as the growth of the human population.<ref name="MJ" /> The production of phosphorus may have peaked before 2011 and some scientists predict reserves will be depleted before the end of the 21st century.<ref name=Lewis /><ref name="MJ" /> Phosphorus comprises about 0.1% by mass of the average rock, and consequently, the Earth's supply is vast, though dilute.<ref name="Greenwood" />

===Wet process===
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 neutralized to give various phosphate salts, which comprise fertilizers. In the wet process, phosphorus does not undergo redox.<ref name=KC/> 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=Taylor>{{cite journal|doi=10.1016/j.jenvman.2009.03.007|pmid=19406560|title=Environmental Impact and Management of Phosphogypsum|journal=Journal of Environmental Management|volume=90|pages=2377–2386|year=2009|last1=Tayibi|first1= Hanan|last2=Choura|first2=Mohamed|last3=López|first3=Félix A.|last4=Alguacil|first4=Francisco J.|last5=López-Delgado|first5=Aurora|issue=8|bibcode=2009JEnvM..90.2377T |hdl=10261/45241|hdl-access=free}}</ref>

===Thermal process===
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=KC>{{cite journal |doi=10.1021/acscentsci.0c00332|title=Let's Make White Phosphorus Obsolete|year=2020|last1=Geeson|first1=Michael B.|last2=Cummins|first2=Christopher C.|journal=ACS Central Science|volume=6|issue=6|pages=848–860|pmid=32607432|pmc=7318074}}</ref> 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 {{chem|SiO|2}}, and [[Coke (fuel)|coke]] to produce {{chem|P|4}}. The {{chem|P|4}} product, being volatile, is readily isolated:<ref>Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company, New York; 2010; p. 379.</ref>

: {{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>{{cite web| url = http://www.heritage.nf.ca/law/erco.html| access-date = 2009-06-06| title= ERCO and Long Harbour| publisher = Memorial University of Newfoundland and the C.R.B. Foundation}}</ref>

===Historical routes===
Historically, before the development of mineral-based extractions, white phosphorus was isolated on an industrial scale from [[bone ash]].<ref>{{cite book|last=Von Wagner|first=Rudolf|title=Manual of chemical technology|date=1897|publisher=D. Appleton & Co.|location=New York|page=411|url=http://babel.hathitrust.org/cgi/pt?id=uc2.ark:/13960/t3tt4gz1p;view=1up;seq=439}}</ref> In this process, the [[tricalcium phosphate]] in bone ash is converted to [[monocalcium phosphate]] with [[sulfuric acid]]:

:{{chem2|Ca3(PO4)2 + 2 H2SO4 -> Ca(H2PO4)2 + 2 CaSO4}}

Monocalcium phosphate is then dehydrated to the corresponding metaphosphate:

:{{chem2|Ca(H2PO4)2 -> Ca(PO3)2 + 2 H2O}}

When ignited to a white heat (~1300 °C) with [[charcoal]], calcium metaphosphate yields two-thirds of its weight of white phosphorus while one-third of the phosphorus remains in the residue as calcium orthophosphate:

:{{chem2|3 Ca(PO3)2 + 10 C -> Ca3(PO4)2 + 10 CO + P4}}

=== Peak phosphorus ===

[[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>{{Cite web |title=Phosphate Rock Statistics and Information {{!}} U.S. Geological Survey |url=https://www.usgs.gov/centers/national-minerals-information-center/phosphate-rock-statistics-and-information |access-date=2023-04-09 |website=www.usgs.gov}}</ref>]]
Peak phosphorus is a concept to describe the point in time when humanity reaches the maximum global production rate of phosphorus as an industrial and commercial [[raw material]]. The term is used in an equivalent way to the better-known term [[peak oil]].<ref name="Cordell2009" /> The issue was raised as a debate on whether phosphorus shortages might be imminent around 2010, which was largely dismissed after [[USGS]] and other organizations<ref name="VanKauwenbergh" /> increased world estimates on available phosphorus resources, mostly in the form of additional resources in [[Morocco]]. However, exact reserve quantities remain uncertain, as do the possible impacts of increased phosphate use on future generations.<ref name="Edixhoven2013">{{cite journal|last1=Edixhoven|first1=J.D.|last2=Gupta|first2=J.|last3=Savenije|first3=H.H.G.|title=Recent revisions of phosphate rock reserves and resources: reassuring or misleading? An in-depth literature review of global estimates of phosphate rock reserves and resources|journal=Earth System Dynamics|date=2013|volume=5|issue=2|pages=491–507|doi=10.5194/esd-5-491-2014|bibcode=2014ESD.....5..491E|doi-access=free}}</ref> This is important because [[Phosphorite|rock phosphate]] is a key ingredient in many inorganic [[Fertilizer|fertilizers]]. Hence, a shortage in rock phosphate (or just significant price increases) might negatively affect the world's [[food security]].<ref name="AmundsonBerhe2015" /> 

Phosphorus is a finite (limited) resource that is widespread in the Earth's crust and in living organisms but is relatively [[scarcity|scarce]] in concentrated forms, which are not evenly distributed across the Earth. The only cost-effective production method to date is the [[mining]] of [[phosphorite|phosphate rock]], but only a few countries have significant commercial [[Mineral resource classification#Mineral reserves and ore reserves|reserves]]. The top five are [[Morocco]] (including reserves located in [[Western Sahara]]), [[China]], [[Egypt]], [[Algeria]] and [[Syria]].<ref name="USGS-2023">{{Cite web |date=January 2023 |title=USGS, Phosphate Rock Statistics and Information |url=https://www.usgs.gov/centers/national-minerals-information-center/phosphate-rock-statistics-and-information |access-date=9 January 2023 |website=Phosphate Rock Statistics and Information}}</ref> 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">{{cite journal|last1=Walan|first1=P.|last2=Davidsson|first2=S.|last3=Johansson|first3=S.|last4=Höök|first4=M.|title=Phosphate rock production and depletion: Regional disaggregated modeling and global implications|journal=Resources, Conservation and Recycling|date=2014|volume=93|issue=12|pages=178–187|doi=10.1016/j.resconrec.2014.10.011|bibcode=2014RCR....93..178W |url=http://www.diva-portal.org/smash/record.jsf?pid=diva2:770437|access-date=9 October 2017}}</ref>

Means of commercial phosphorus production besides mining are few because the [[phosphorus cycle]] does not include significant gas-phase transport.<ref name="Neset">{{cite journal|title=Global phosphorus scarcity: identifying synergies for a sustainable future|journal=Journal of the Science of Food and Agriculture|last1=Neset|first1=Tina-Simone S.|last2=Cordell|first2=Dana|volume=92|issue=1|pages=2–6|year=2011|doi=10.1002/jsfa.4650|pmid=21969145}}</ref> The predominant source of phosphorus in modern times is phosphate rock (as opposed to the guano that preceded it). According to some researchers, Earth's commercial and affordable phosphorus reserves are expected to be depleted in 50–100 years and peak phosphorus to be reached in approximately 2030.<ref name="Cordell2009">{{harvnb|Cordell|Drangert|White|2009}}</ref><ref name="Lewis">{{cite news|title=Scientists warn of lack of vital phosphorus as biofuels raise demands|newspaper=[[The Times]]|last1=Lewis|first1=Leo|date=23 June 2008|url=http://business.timesonline.co.uk/tol/business/industry_sectors/natural_resources/article4193017.ece|url-status=dead|archive-url=https://web.archive.org/web/20110723130701/http://www.foodandwatersecurity.net/data/172.pdf|archive-date=23 July 2011}}</ref> Others suggest that supplies will last for several hundreds of years.<ref name="ifdc.org">{{cite web |title=IFDC Report Indicates Adequate Phosphorus Resources Available to Meet Global Food Demands |url=https://ifdc.org/2010/09/22/ifdc-report-indicates-adequate-phosphorus-resources-available-to-meet-global-food-demands/ |date=22 September 2010}}</ref> As with the [[predicting the timing of peak oil|timing of peak oil]], the question is not settled, and researchers in different fields regularly publish different estimates of the rock phosphate reserves.<ref name="Edixhoven">{{cite journal|last1=Edixhoven|first1=J. D.|last2=Gupta|first2=J.|last3=Savenije|first3=H. H. G.|title=Recent revisions of phosphate rock reserves and resources: a critique|journal=Earth System Dynamics|volume=5|issue=2|year=2014|pages=491–507|issn=2190-4987|doi=10.5194/esd-5-491-2014|bibcode=2014ESD.....5..491E|url=https://pure.uva.nl/ws/files/2486540/162701_478483.pdf|doi-access=free}}</ref>

====Background====
[[File:US Mined Phosphate Rock 1900-2015.png|thumb|upright=1.5|Phosphate rock mined in the United States, 1900–2015 (data from US Geological Survey)]]

The peak phosphorus concept is connected with the concept of [[planetary boundaries]]. Phosphorus, as part of [[biogeochemistry|biogeochemical]] processes, belongs to one of the nine "Earth system processes" which are known to have boundaries. As long as the boundaries are not crossed, they mark the "safe zone" for the planet.<ref>{{cite journal | last1 = Rockström | first1 = J. | last2 = Steffen | first2 = K. |display-authors=etal  | year = 2009 | title = Planetary boundaries: exploring the safe operating space for humanity | url = http://www.stockholmresilience.org/download/18.8615c78125078c8d3380002197/ES-2009-3180.pdf | journal = Ecology and Society | volume = 14 | issue = 2| page = 32 | doi = 10.5751/ES-03180-140232 | doi-access = free }}</ref>

====Estimates of world phosphate reserves====
[[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<ref>Arno Rosemarin (2016) [https://dakofa.com/fileadmin/user_upload/1600_Arno_Rosemarin_Stockholm_Environment_Institute.pdf Phosphorus a Limited Resource – Closing the Loop], Global Status of Phosphorus Conference, Malmö, Sweden (based on [https://minerals.usgs.gov/minerals/pubs/commodity/phosphate_rock/ USGS Phosphate Rock Statistics and Information])</ref>]]

The accurate determination of peak [[phosphorus]] is dependent on knowing the total world's commercial [[phosphate]] reserves and resources, especially in the form of [[phosphate rock]] (a summarizing term for 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="Sutton" /><ref>{{Cite book|title=CIM DEFINITION STANDARDS - For Mineral Resources and Mineral Reserves|publisher=CIM Standing Committee on Reserve Definitions|year=2010|url=http://web.cim.org/userfiles/file/cim_definiton_standards_nov_2010.pdf|pages=4–6|archive-url=https://web.archive.org/web/20190214115454/http://web.cim.org/userfiles/file/cim_definiton_standards_nov_2010.pdf|archive-date=14 February 2019|accessdate=5 April 2024|url-status=dead}}</ref>

Unprocessed phosphate rock has a concentration of 1.7–8.7% phosphorus by mass (4–20% [[phosphorus pentoxide]]). By comparison, the Earth's crust contains 0.1% phosphorus by mass,<ref>U.S. Geological Survey [http://pubs.usgs.gov/of/2004/1368/Soil_PDFs/P_soils_page.pdf Phosphorus Soil Samples]</ref> and vegetation 0.03–0.2%.<ref>[http://www.seafriends.org.nz/oceano/abund.htm Abundance of Elements]</ref> Although quadrillions of tons of phosphorus exist in the Earth's crust,<ref>American Geophysical Union, Fall Meeting 2007, abstract #V33A-1161. [http://adsabs.harvard.edu/abs/2007AGUFM.V33A1161P Mass and Composition of the Continental Crust]</ref> these are currently not economically extractable.

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="USGS-2023" /> 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="ifdc.org"/><ref name="VanKauwenbergh">{{cite book|last1=Van Kauwenbergh|first1=Steven J.|title=World Phosphate Rock Reserves and Resources|date=2010|publisher=[[International Fertilizer Development Center]] (IFDC)|location=Muscle Shoals, AL, USA|isbn=978-0-88090-167-3|pages=60|url=http://ifdc.org/technical-bulletins/|access-date=7 April 2016|archive-date=19 August 2018|archive-url=https://web.archive.org/web/20180819051418/https://ifdc.org/technical-bulletins/|url-status=dead}}</ref> Phosphorus reserve figures are intensely debated.<ref name="Sutton">{{cite book|title=Our Nutrient World: The challenge to produce more food and energy with less pollution|url=http://www.initrogen.org/sites/default/files/documents/files/ONW.pdf|last1=Sutton|first1=M.A.|last2=Bleeker, A.|last3=Howard, C.M.|publisher=Centre for Ecology and Hydrology, Edinburgh on behalf of the Global Partnership on Nutrient Management and the International Nitrogen Initiative.|date=2013|isbn=978-1-906698-40-9|display-authors=etal|access-date=2015-05-12|archive-url=https://web.archive.org/web/20161104175311/http://www.initrogen.org/sites/default/files/documents/files/ONW.pdf|archive-date=2016-11-04|url-status=dead}}</ref>{{sfn|Cordell|White|2011}}<ref name="Van VuurenBouwman2010">{{cite journal|last1=Van Vuuren|first1=D.P.|last2=Bouwman|first2=A.F.|last3=Beusen|first3=A.H.W.|title=Phosphorus demand for the 1970–2100 period: A scenario analysis of resource depletion|journal=Global Environmental Change|volume=20|issue=3|year=2010|pages=428–439|issn=0959-3780|doi=10.1016/j.gloenvcha.2010.04.004|bibcode=2010GEC....20..428V }}</ref> Gilbert suggest that there has been little external verification of the estimate.<ref name="Gilbert">{{cite journal|last=Gilbert|first=Natasha|title=The disappearing nutrient|journal=Nature|volume=461|issue=7265|pages=716–718|doi=10.1038/461716a|pmid=19812648|date=8 October 2009|doi-access=|s2cid=4419892 }}</ref> A 2014 review<ref name="Edixhoven" /> 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".

The countries with most phosphate rock commercial reserves (in billion metric tons): [[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, [[Finland]] 1.0, [[Russia]] 0.6, [[Jordan]] 0.8.<ref>{{Cite web|url=http://verkkolehti.geofoorumi.fi/en/2015/10/finlands-phosphorus-resources-are-more-important-than-ever/|title=Finland's phosphorus resources are more important than ever (Geological Survey of Finland)|last=Ahokas|first=K.|date=2015|access-date=2017-04-01|archive-date=2019-05-06|archive-url=https://web.archive.org/web/20190506172330/http://verkkolehti.geofoorumi.fi/en/2015/10/finlands-phosphorus-resources-are-more-important-than-ever/|url-status=dead}}</ref><ref name="USGS-2023" />

Rock phosphate shortages (or just significant price increases) might negatively affect the world's [[food security]].<ref name="AmundsonBerhe2015">{{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.|title=Soil and human security in the 21st century|journal=Science|volume=348|issue=6235|year=2015|pages=1261071|issn=0036-8075|doi=10.1126/science.1261071|pmid=25954014|bibcode=2015Sci...34861071A |s2cid=206562728|url=http://www.escholarship.org/uc/item/8f42m6w4}}</ref> Many agricultural systems depend on supplies of inorganic fertilizer, which use rock phosphate. Under the food production regime in developed countries, shortages of rock phosphate could lead to shortages of inorganic fertilizer, which could in turn reduce the global food production.<ref name="Pollan">{{cite book|title=The Omnivore's Dilemma: A Natural History of Four Meals|last=Pollan|first=Michael|isbn=978-1-59420-082-3|date=11 April 2006|publisher=Penguin Press|url-access=registration|url=https://archive.org/details/omnivoresdilemma00poll_0}}</ref>

Economists have pointed out that price fluctuations of rock phosphate do not necessarily indicate peak phosphorus, as these have already occurred due to various demand- and supply-side factors.{{sfn|Heckenmüller|Narita|Klepper|2014}}

=====United States=====
{{Further|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>US Geological Survey, [https://pubs.usgs.gov/periodicals/mcs2021/mcs2021-phosphate.pdf Phosphate Rock], 2021.</ref>

=====Exhaustion of guano reserves=====
{{Main|Guano}}
In 1609 [[Garcilaso de la Vega (El 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 fertilizer. As Garcilaso described, the Incas near the coast harvested guano.<ref name="Leigh">{{cite book|title=The World's Greatest Fix: A History of Nitrogen and Agriculture|url=https://archive.org/details/worldsgreatestfi0000leig|url-access=registration|last=Leigh|first=G. J.|year=2004|isbn=978-0-19-516582-1|publisher=Oxford University Press}}</ref> In the early 1800s [[Alexander von Humboldt]] introduced [[guano]] as a source of [[agriculture|agricultural]] [[fertilizer]] to Europe after having discovered it 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 {{nowrap|30 meters}} deep.<ref name="Skaggs">{{cite book|title=The Great Guano Rush: Entrepreneurs and American Overseas Expansion|last=Skaggs|first=Jimmy M.|publisher=St. Martin's Press|date=May 1995|isbn=978-0-312-12339-0}}</ref> The guano had previously been used by the [[Moche culture|Moche]] people as a source of fertilizer by mining it and transporting it back to [[Peru]] by boat. International commerce in guano did not start until after 1840.<ref name="Skaggs" /> 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]].

====Phosphorus conservation and recycling====
[[File:Nauru-phosphatefields.jpg|thumb|Phosphate mine on [[Nauru]], once one of the world's major sources of phosphate rock]]

===== Overview =====
Phosphorus can be transferred from the soil in one location to another as food is transported across the world, taking the phosphorus it contains with it. Once consumed by humans, it can end up in the local environment (in the case of [[open defecation]] which is still widespread on a global scale) or in rivers or the ocean via [[sanitary sewer|sewage systems]] and [[sewage treatment plant]]s in the case of cities connected to sewer systems. An example of one crop that takes up large amounts of phosphorus is [[soy]].

In an effort to postpone the onset of peak phosphorus several methods of reducing and reusing phosphorus are in practice, such as in agriculture and in [[sanitation]] systems. The [[Soil Association]], the UK organic agriculture certification and pressure group, issued a report in 2010 "A Rock and a Hard Place" encouraging more recycling of phosphorus.<ref>[http://www.soilassociation.org/LinkClick.aspx?fileticket=eeGPQJORrkw%3D&tabid=57 soilassociation.org - A rock and a hard place, Peak phosphorus and the threat to our food security] {{webarchive|url=https://web.archive.org/web/20101223135324/http://soilassociation.org/LinkClick.aspx?fileticket=eeGPQJORrkw%3D&tabid=57 |date=2010-12-23 }}, 2010</ref> One potential solution to the shortage of phosphorus is greater recycling of human and animal wastes back into the environment.<ref name="Miller-McCune">{{harvnb|Burns|2010}}</ref>

===== Agricultural practices =====
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. These methods are still dependent on a periodic application of phosphate rock to the soil and as such methods to recycle the lost phosphorus have also been proposed. Perennial vegetation, such as grassland or forest, is much more efficient in its use of phosphate than arable land. Strips of grassland and/or forest between arable land and rivers can greatly reduce losses of phosphate and other nutrients.<ref>{{cite journal|title=Phosphorus and nitrogen losses in relation to forest, pasture and row-crop land use and precipitation distribution in the midwest usa|journal=Journal of Water Science|last1=Udawatta|first1=Ranjith P.|last2=Henderson|first2=Gray S.|last3=Jones|first3=John R.|last4=Hammer|first4=David
|volume=24|issue=3|pages=269–281|year=2011|doi=10.7202/1006477ar|doi-access=free}}</ref>

Integrated farming systems which use animal sources to supply phosphorus for crops do exist at smaller scales, and application of the system to a larger scale is a potential alternative for supplying the nutrient, although it would require significant changes to the widely adopted modern crop fertilizing methods.

===== Excreta reuse =====
{{Main|Excreta reuse}}
The oldest method of recycling phosphorus is through the reuse of animal [[manure]] and human [[excreta]] in agriculture. Via this method, phosphorus in the foods consumed are excreted, and the animal or human excreta are subsequently collected and re-applied to the fields. Although this method has maintained civilizations for centuries the current system of manure management is not logistically geared towards application to crop fields on a large scale. At present, manure application could not meet the phosphorus needs of large scale agriculture. Despite that, it is still an efficient method of recycling used phosphorus and returning it to the soil. There are concerns with pathogens in manure and human excreta, but those pathogens can be eliminated via suitable treatment. However, especially in the [[Global South]] these processes are not always followed, leading to outbreaks of diseases transmitted via the [[fecal–oral route]] such as [[cholera]].

===== Sewage sludge =====
Sewage treatment plants that have an [[enhanced biological phosphorus removal]] step produce a [[sewage sludge]] that is rich in phosphorus. Various processes have been developed to extract phosphorus from sewage sludge directly, from the ash after [[incineration]] of the sewage sludge or from other products of [[sewage sludge treatment]]. This includes the extraction of phosphorus rich materials such as [[struvite]] from waste processing plants.<ref name="Gilbert" /> The struvite can be made by adding magnesium to the waste. Some companies such as Ostara in Canada and NuReSys in Belgium are already using this technique to recover phosphate.<ref>{{Cite web |last=Tweed |first=Katherine |date=2009-11-01 |title=Sewage Industry Fights Phosphorus Pollution |url=https://www.scientificamerican.com/article/sewages-cash-crop/ |access-date=2024-06-21 |website=Scientific American |language=en}}</ref>

Research on phosphorus recovery methods from sewage sludge has been carried out in Sweden and Germany since around 2003, but the technologies currently under development are not yet cost effective, given the current price of phosphorus on the world market.<ref>Sartorius, C., von Horn, J., Tettenborn, F. (2011). [http://www.susana.org/en/resources/library/details/1304 Phosphorus recovery from wastewater – state-of-the-art and future potential]. Conference presentation at Nutrient Recovery and Management Conference organised by International Water Association (IWA) and Water Environment Federation (WEF) in Florida, USA</ref><ref>Hultman, B., Levlin, E., Plaza, E., Stark, K. (2003). [http://www.susana.org/en/resources/library/details/436 Phosphorus Recovery from Sludge in Sweden - Possibilities to meet proposed goals in an efficient, sustainable and economical way].</ref>
===Neutron transmutation doping===
The above routes refer to "production" in the ''chemical'' sense i.e. extracting a desired element or compound from a source without changing the atoms themselves. However, there is a process which produces phosphorus in a ''nuclear'' sense in that atoms of another element are turned into phosphorus. While the amount of phosphorus produced this way is minuscule, it is nonetheless a crucial process in semiconductor production.

Neutron [[nuclear transmutation|transmutation]] doping (NTD) is an unusual doping method for special applications. Most commonly, it is used to dope silicon n-type in high-power electronics and [[semiconductor detector]]s. It is based on the conversion of the <sup>30</sup>Si isotope into phosphorus atoms by  [[neutron absorption]] 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]] (most commonly a [[research reactor]] e.g. the one at [[Massachusetts Institute of Technology|MIT]]<ref>{{cite web | url=http://nrl.mit.edu/facilities/ntds | title=Neutron Transmutation Doping of Silicon &#124; MIT Nuclear Reactor Laboratory }}</ref>) to receive the neutrons. As neutrons continue to pass through the silicon, more and more phosphorus atoms are produced by transmutation, and therefore the doping becomes more and more strongly n-type. NTD is a far less common doping method than diffusion or ion implantation, but it has the advantage of creating an extremely uniform dopant distribution.<ref>{{cite book | last=Baliga | first=B. Jayant | title=Modern Power Devices | publisher=Wiley-Interscience | date=1987-03-10 | isbn=0-471-81986-7 |page=32}}</ref><ref>{{cite conference|last1=Schmidt |first1=P. E. |last2=Vedde |first2=J. |year=1998| title=High Resistivity NTD Production and Applications |conference=Electrochemical Society Proceedings|volume=98 |issue=13|isbn=9781566772075}}</ref>

==Applications==

===Flame retardant===
Phosphorus compounds are used as flame retardants. Flame-retardant materials and coatings are being developed that are both phosphorus- and bio-based.<ref>{{Cite journal |last1=Naiker |first1=Vidhukrishnan E. |last2=Mestry |first2=Siddhesh |last3=Nirgude |first3=Tejal |last4=Gadgeel |first4=Arjit |last5=Mhaske |first5=S. T. |date=2023-01-01 |title=Recent developments in phosphorous-containing bio-based flame-retardant (FR) materials for coatings: an attentive review |journal=Journal of Coatings Technology and Research |language=en |volume=20 |issue=1 |pages=113–139 |doi=10.1007/s11998-022-00685-z |s2cid=253349703 |issn=1935-3804}}</ref>

===Food additive===
Phosphorus is an essential [[Mineral (nutrient)|mineral]] for humans listed in the [[Dietary Reference Intake#Minerals|Dietary Reference Intake]] (DRI).

Food-grade [[phosphoric acid]] (additive [[E number|E338]]<ref name=fgovuk>{{cite web|url=http://www.food.gov.uk/policy-advice/additivesbranch/enumberlist#h_7|title=Current EU approved additives and their E Numbers|date=14 March 2012|publisher=Foods Standards Agency|access-date=22 July 2012|archive-date=21 August 2013|archive-url=https://web.archive.org/web/20130821045312/http://food.gov.uk/policy-advice/additivesbranch/enumberlist#h_7|url-status=live}}</ref>) is used to acidify foods and beverages such as various [[cola]]s and jams, providing a tangy or sour taste. The phosphoric acid also serves as a [[preservative]].<ref>{{Cite web|title=Why is phosphoric acid used in some Coca‑Cola drinks?{{!}} Frequently Asked Questions {{!}} Coca-Cola GB|url=https://www.coca-cola.co.uk/our-business/faqs/why-is-phosphoric-acid-used-in-coca-cola-drinks-diet-coke-coke-zero|access-date=2021-08-31|website=www.coca-cola.co.uk|language=en-GB|archive-date=2 August 2021|archive-url=https://web.archive.org/web/20210802114054/https://www.coca-cola.co.uk/our-business/faqs/why-is-phosphoric-acid-used-in-coca-cola-drinks-diet-coke-coke-zero|url-status=live}}</ref> 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>{{Cite journal|title=Dietary advice in dental practice|journal=British Dental Journal|volume=193|issue=10|pages=563–568|date=23 November 2002|doi=10.1038/sj.bdj.4801628|pmid=12481178|last1=Moynihan|first1=P. J.|doi-access=free}}</ref> Phosphoric acid also has the potential to contribute to the formation of [[Kidney stone disease|kidney stones]], especially in those who have had kidney stones previously.<ref name=Qa2014>{{cite journal |last1= Qaseem |first1= A |last2= Dallas |first2= P |last3= Forciea |first3= MA |last4= Starkey |first4= M |last5= Denberg |first5= TD |display-authors= 4 |title= Dietary and pharmacologic management to prevent recurrent nephrolithiasis in adults: A clinical practice guideline from the American College of Physicians |journal= [[Annals of Internal Medicine]] |date= 4 November 2014 |volume= 161 |issue= 9 |pages= 659–67 |doi= 10.7326/M13-2908 |pmid=25364887|doi-access= free }}</ref>

===Fertiliser===
{{Main|Fertiliser}}
{{See also|Phosphorus cycle}}
Phosphorus is an essential plant nutrient (the most often limiting nutrient, after [[nitrogen]]),<ref name="et">{{cite book |last=Etesami |first = H. |title=Nutrient Dynamics for Sustainable Crop Production |date=2019 |page=217 |publisher = Springer |isbn = 978-981-13-8660-2 |url=https://books.google.com/books?id=DeKtDwAAQBAJ&q=phosphorous%20limiting}}</ref> and the bulk of all phosphorus production is in concentrated phosphoric acids for [[agriculture]] [[fertiliser]]s, containing as much as 70% to 75% P<sub>2</sub>O<sub>5</sub>. That led to large increase in [[phosphate]] (PO<sub>4</sub><sup>3−</sup>) production in the second half of the 20th century.<ref name="MJ">{{cite magazine|url=https://www.motherjones.com/environment/2013/05/fertilizer-peak-phosphorus-shortage|title=You Need Phosphorus to Live—and We're Running Out|last=Philpott|first=Tom|date=March–April 2013|magazine=Mother Jones}}</ref> 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="et" /> Heavy use of phosphorus fertilizers and their runoff have resulted in [[eutrophication]] (overenrichment) of [[aquatic ecosystem]]s.<ref name="Carpenter 2005">{{cite journal | last=Carpenter | first=Stephen R. | title=Eutrophication of aquatic ecosystems: Bistability and soil phosphorus | journal=Proceedings of the National Academy of Sciences | volume=102 | date=2005 | issue=29 | issn=0027-8424 | doi=10.1073/pnas.0503959102 | pages=10002–10005| pmid=15972805 | pmc=1177388 | doi-access=free | bibcode=2005PNAS..10210002C }}</ref><ref name="Conley 2009">{{cite journal | last1=Conley | first1=Daniel J. | last2=Paerl | first2=Hans W. | last3=Howarth | first3=Robert W. | display-authors=etal | title=Controlling Eutrophication: Nitrogen and Phosphorus | journal=Science | volume=323 | date=2009 | issue=5917 | issn=0036-8075 | doi=10.1126/science.1167755 | pages=1014–1015| pmid=19229022 }}</ref>

Natural phosphorus-bearing compounds are mostly inaccessible to plants because of the low solubility and mobility in soil.<ref>{{cite web |title=Soil Phosphorous |url=https://www.nrcs.usda.gov/Internet/FSE_DOCUMENTS/nrcs142p2_053254.pdf |website=United States Department of Agriculture |access-date=2020-08-17 |archive-date=2020-10-28 |archive-url=https://web.archive.org/web/20201028202404/https://www.nrcs.usda.gov/Internet/FSE_DOCUMENTS/nrcs142p2_053254.pdf |url-status=dead }}</ref> Most phosphorus is very stable in the soil minerals or organic matter of the soil. Even when phosphorus is added in manure or fertilizer 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>{{cite web |title= Managing Phosphorus for Crop Production |url= https://extension.psu.edu/programs/nutrient-management/educational/soil-fertility/managing-phosphorus-for-crop-production#:~:text=The%20challenge%20is%20that%20phosphorus,only%20from%20the%20soil%20solution |website= Penn State Extension |access-date= 2020-08-17 |archive-date= 2020-10-20 |archive-url= https://web.archive.org/web/20201020090515/https://extension.psu.edu/programs/nutrient-management/educational/soil-fertility/managing-phosphorus-for-crop-production#:~:text=The%20challenge%20is%20that%20phosphorus,only%20from%20the%20soil%20solution |url-status= dead }}</ref> Fertiliser is often in the form of superphosphate of lime, a mixture of calcium dihydrogen phosphate (Ca(H<sub>2</sub>PO<sub>4</sub>)<sub>2</sub>), and calcium sulfate dihydrate (CaSO<sub>4</sub>·2H<sub>2</sub>O) 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>{{cite book|title=Industrial Minerals & Rocks: Commodities, Markets, and Uses|editor=Jessica Elzea Kogel|publisher=SME, 2006|isbn=0-87335-233-5|page=964|year=2006}}</ref>

{|class="wikitable"
|-
!Widely used compounds!!Use
|-
|[[monocalcium phosphate|Ca(H<sub>2</sub>PO<sub>4</sub>)<sub>2</sub>·H<sub>2</sub>O]]||Baking powder and fertilisers
|-
|[[dicalcium phosphate|CaHPO<sub>4</sub>·2H<sub>2</sub>O]]||Animal food additive, toothpowder
|-
|[[phosphoric acid|H<sub>3</sub>PO<sub>4</sub>]]||Manufacture of phosphate fertilisers
|-
|[[phosphorus trichloride|PCl<sub>3</sub>]]||Manufacture of POCl<sub>3</sub> and pesticides
|-
|[[phosphoryl chloride|POCl<sub>3</sub>]]||Manufacture of plasticiser
|-
|[[phosphorus pentasulfide|P<sub>4</sub>S<sub>10</sub>]]||Manufacturing of additives and pesticides
|-
|[[sodium triphosphate|Na<sub>5</sub>P<sub>3</sub>O<sub>10</sub>]]||Detergents
|}

===Organophosphorus===
White phosphorus is widely used to make [[organophosphorus compound]]s through intermediate [[phosphorus chlorides]] and two phosphorus sulfides, [[phosphorus pentasulfide]] and [[phosphorus sesquisulfide]].{{sfn|Threlfall|1951}} Organophosphorus compounds have many applications, including in [[plasticizer|plasticisers]], [[flame retardant]]s, [[pesticide]]s, extraction agents, nerve agents and [[water treatment]].<ref name="Greenwood" />{{sfn|Diskowski|Hofmann}}

===Metallurgical aspects===
Phosphorus is also an important component in [[steel]] production, in the making of [[phosphor bronze]], and in many other related products.<ref>{{cite book|title=Sustainable Phosphorus Management: A Global Transdisciplinary Roadmap|publisher=Springer Science & Business Media |editor=Roland W. Scholz |editor2=Amit H. Roy |editor3=Fridolin S. Brand |editor4=Deborah Hellums |editor5=Andrea E. Ulrich|isbn=978-94-007-7250-2|page=175|date=2014-03-12}}</ref><ref>{{cite book|title=Encyclopedia and Handbook of Materials, Parts and Finishes|publisher=CRC Press |author=Mel Schwartz|isbn=978-1-138-03206-4|date=2016-07-06}}</ref> 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>{{cite book|title=Copper and Copper Alloys|publisher=ASM International |editor=Joseph R. Davisz|isbn=0-87170-726-8|page=181|date=January 2001}}</ref>
[[Phosphate conversion coating]] is a chemical treatment applied to steel parts to improve their corrosion resistance.

===Matches===
[[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.]]
{{main|Match}}
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,<ref>{{cite journal|journal=Br. J. Ind. Med.|year=1962|volume=19|pages=83–99|title=Phosphorus Necrosis of the Jaw: A Present-day Study: With Clinical and Biochemical Studies|author=Hughes, J. P. W |author2=Baron, R. |author3=Buckland, D. H. |author4=Cooke, M. A. |author5=Craig, J. D. |author6=Duffield, D. P. |author7=Grosart, A. W. |author8=Parkes, P. W. J. |author9=Porter, A. |display-authors=3 |pmc=1038164|issue=2|pmid=14449812|doi=10.1136/oem.19.2.83}}</ref> sensitive to storage conditions, toxic if ingested, and hazardous when accidentally ignited on a rough surface.<ref name="crass9">{{cite journal|title=A history of the match industry. Part 9| author=Crass, M. F. Jr. |year=1941|pages=428–431|journal=Journal of Chemical Education|volume=18|url=http://www.jce.divched.org/journal/Issues/1941/Sep/jceSubscriber/JCE1941p0428.pdf|bibcode=1941JChEd..18..428C|doi=10.1021/ed018p428|issue=9}}{{dead link|date=March 2018 |bot=InternetArchiveBot |fix-attempted=yes }}</ref><ref>{{cite journal|title=Industrial disease due to certain poisonous fumes or gases|author=Oliver, Thomas|url=https://archive.org/stream/archivesofpublic01victuoft#page/2/mode/1up|pages=1–21|journal=Archives of the Public Health Laboratory|volume=1|publisher=Manchester University Press|year=1906}}</ref> Production in several countries was banned between 1872 and 1925.<ref>{{cite journal|first=Steve|last= Charnovitz |title=The Influence of International Labour Standards on the World Trading Regime. A Historical Overview | journal= International Labour Review| volume= 126| issue= 5| date= 1987| pages=565, 571}}</ref> The international [[Berne Convention (1906)|Berne Convention]], ratified in 1906, prohibited the use of white phosphorus in matches.

In consequence, phosphorous matches were gradually replaced by safer alternatives. 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]] (P<sub>4</sub>S<sub>3</sub>), 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 safety match.

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. Nonetheless, safety matches, invented in 1844 by [[Gustaf Erik Pasch]] and market ready by the 1860s, did not gain consumer acceptance until the prohibition of white phosphorus. Using a dedicated striker strip was considered clumsy.<ref name="wiberg" />{{sfn|Threlfall|1951}}<ref name="Hardt_2001">{{Cite book|author=Alexander P. Hardt|title=Pyrotechnics|publisher=Pyrotechnica Publications|location=Post Falls Idaho US|date=2001|isbn=0-929388-06-2|chapter=Matches|pages=74–84}}</ref>

===Water softening===
[[Sodium tripolyphosphate]] made from phosphoric acid is used in laundry detergents in some countries, but banned for this use in others.<ref name="CRC" /> This compound softens the water to enhance the performance of the detergents and to prevent pipe/boiler tube [[corrosion]].{{sfn|Schrödter|Bettermann|Staffel|Wahl}}

===Miscellaneous===
* Phosphates are used to make special glasses for [[sodium lamp]]s.<ref name="CRC" />
* Bone-ash (mostly [[calcium phosphate]]) is used in the production of fine china.<ref name="CRC">{{Cite book| author = Hammond, C. R. |chapter = The Elements |title=Handbook of Chemistry and Physics |edition = 81st| publisher =CRC press| date = 2000| isbn = 0-8493-0481-4}}</ref>
* Phosphoric acid made from elemental phosphorus is used in food applications such as [[Soft drink#Phosphate soda|soft drinks]], and as a starting point for food grade phosphates.{{sfn|Threlfall|1951}} These include [[monocalcium phosphate]] for [[baking powder]] and [[sodium tripolyphosphate]].{{sfn|Threlfall|1951}} Phosphates are used to improve the characteristics of processed meat and [[cheese]], and in toothpaste.{{sfn|Threlfall|1951}}
* [[White phosphorus munitions|White phosphorus]], called "WP" (slang term "Willie Peter") is used in [[military]] applications as [[incendiary device|incendiary bombs]], for [[smoke-screen]]ing as smoke pots and [[smoke bomb]]s, and in [[tracer ammunition]]. It is also a part of an obsolete [[M34 grenade|M34 White Phosphorus US hand grenade]]. This multipurpose grenade was mostly used for signaling, smoke screens, and inflammation; it could also cause severe burns and had a psychological impact on the enemy.<ref>{{Cite book| author = Dockery, Kevin|title = Special Warfare Special Weapons|location = Chicago|publisher = Emperor's Press| date = 1997|isbn = 1-883476-00-3}}</ref> Military uses of white phosphorus are constrained by international law.
* <sup>32</sup>P and <sup>33</sup>P are used as radioactive tracers in biochemical laboratories.<ref>{{cite book|title=Radionuclides in the Environment|editor=David A. Atwood|publisher=John Wiley & Sons, 2013|isbn=978-1-118-63269-7|date=2013-02-19}}</ref>
* Phosphorus is a [[doping (semiconductor)|dopant]] in [[n-type semiconductor]]s

==Biological role==
{{see also|Calcium metabolism}}
Inorganic phosphorus in the form of the phosphate {{chem|PO|4|3-}} is required for all known forms of [[life]].{{sfn|Ruttenberg}} 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. [[Phospholipid]]s are the main structural components of all cellular membranes. [[Calcium phosphate]] salts assist in stiffening [[bone]]s.<ref name="Greenwood" /> Biochemists commonly use the abbreviation "P{{sub|i}}" to refer to inorganic phosphate.<ref>{{cite journal | last1 = Lipmann | first1 = D. | year = 1944 | title = Enzymatic Synthesis of Acetyl Phosphate | url = http://www.jbc.org/content/155/1/55.short | journal = J Biol Chem | volume = 155 | pages = 55–70 | doi = 10.1016/S0021-9258(18)43172-9 | doi-access = free }}</ref>

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. Phospholipids 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>Nelson, D. L.; Cox, M. M. "Lehninger, Principles of Biochemistry" 3rd Ed. Worth Publishing: New York, 2000. {{ISBN|1-57259-153-6}}.</ref>

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. Average phosphorus concentration in the blood is about 0.4&nbsp;g/L; about 70% of that is organic and 30% inorganic phosphates.<ref>{{Cite book |url=https://books.google.com/books?id=ba_5OSsyS4YC&pg=PA171 |page=171 |title=Nutrition for the Middle Aged and Elderly |author= Bernhardt, Nancy E. |author2= Kasko, Artur M. |publisher=Nova Publishers |date=2008 |isbn=978-1-60456-146-3}}</ref> An adult with healthy diet consumes and excretes about 1–3&nbsp;grams of phosphorus per day, with consumption in the form of inorganic phosphate and phosphorus-containing biomolecules such as [[nucleic acids]] and [[phospholipids]]; and excretion almost exclusively in the form of phosphate ions such as {{chem|H|2|PO|4|-}} and {{chem|HPO|4|2-}}. Only about 0.1% of body phosphate circulates in the blood, paralleling the amount of phosphate available to soft tissue cells.

===Bone and teeth enamel===
The main component of bone is [[hydroxyapatite]] as well as amorphous forms of calcium phosphate, possibly including carbonate. Hydroxyapatite is the main component of tooth enamel. [[Water fluoridation]] enhances the resistance of teeth to decay by the partial conversion of this mineral to the still harder material [[fluorapatite]]:<ref name="Greenwood" />
: {{chem|Ca|5|(|P|O|4|)|3|O|H}} + {{chem|F|-}} → {{chem|Ca|5|(|P|O|4|)|3|F}} + {{chem|O|H|-}}

===Phosphorus deficiency===
In medicine, phosphate deficiency syndrome 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="pmid18583681">{{cite journal |author=Mehanna H. M. |author2=Moledina J. |author3=Travis J. |title=Refeeding syndrome: what it is, and how to prevent and treat it |journal=BMJ |volume=336 |issue=7659 |pages=1495–8 |date=June 2008 |pmid=18583681 |pmc=2440847 |doi=10.1136/bmj.a301 }}</ref>) 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>{{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 |pmid=8642449 |doi=10.1093/jn/126.suppl_4.1153S |doi-access=free }}</ref>

Phosphorus is an essential [[macromineral]] for plants, which is studied extensively in [[edaphology]] to understand plant uptake from [[soil]] systems. Phosphorus is a [[limiting factor]] in many [[ecosystem]]s; that is, the scarcity of phosphorus limits the rate of organism growth. An excess of phosphorus can also be problematic, especially in aquatic systems where [[eutrophication]] sometimes leads to [[algal blooms]].<ref name="MJ" />

=== Nutrition ===
==== Dietary recommendations ====
The [[U.S. Institute of Medicine]] (IOM) updated Estimated Average Requirements (EARs) and [[Recommended Dietary Allowance]]s (RDAs) for phosphorus in 1997. If there is not sufficient information to establish EARs and RDAs, an estimate designated [[Adequate Intake]] (AI) is used instead. The current EAR 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]]s (ULs) for vitamins and minerals when evidence is sufficient. In the case of phosphorus, the UL is 4000&nbsp;mg/day. Collectively the EARs, RDAs, AIs and ULs are referred to as [[Dietary Reference Intake]]s (DRIs).<ref name="DRItext">{{cite book | last1 = Institute of Medicine | title = Dietary Reference Intakes for Calcium, Phosphorus, Magnesium, Vitamin D, and Fluoride | chapter = Phosphorus | publisher = The National Academies Press | year = 1997 | location = Washington, DC | pages = 146–189 | chapter-url = https://www.nap.edu/read/5776/chapter/7| doi = 10.17226/5776 | pmid = 23115811 | isbn = 978-0-309-06403-3 | s2cid = 8768378 | author1-link = Institute of Medicine }}</ref>

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="efsa24">{{cite web |title=Dietary reference values |url=https://www.efsa.europa.eu/en/topics/topic/dietary-reference-values |publisher=European Food Safety Authority |access-date=24 September 2024 |date=7 August 2024}}</ref> 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>{{cite web | title = Overview on Dietary Reference Values for the EU population as derived by the EFSA Panel on Dietetic Products, Nutrition and Allergies | year = 2017 | url = https://www.efsa.europa.eu/sites/default/files/assets/DRV_Summary_tables_jan_17.pdf}}</ref> EFSA reviewed the same safety question and decided that there was not sufficient information to set a UL.<ref>{{citation | title = Tolerable Upper Intake Levels For Vitamins And Minerals | publisher = European Food Safety Authority | year = 2006 | url = http://www.efsa.europa.eu/sites/default/files/efsa_rep/blobserver_assets/ndatolerableuil.pdf}}</ref>

For U.S. food and dietary supplement labeling purposes, the amount in a serving is expressed as a percent of Daily Value (%DV). For phosphorus labeling purposes, 100% of the Daily Value was 1000&nbsp;mg, but as of May&nbsp;27, 2016, it was revised to 1250&nbsp;mg to bring it into agreement with the RDA.<ref name="FedReg">{{Cite web |url=https://www.gpo.gov/fdsys/pkg/FR-2016-05-27/pdf/2016-11867.pdf |title=Federal Register May 27, 2016 Food Labeling: Revision of the Nutrition and Supplement Facts Labels. FR page 33982}}</ref><ref>{{cite web | title=Daily Value Reference of the Dietary Supplement Label Database (DSLD) | website=Dietary Supplement Label Database (DSLD) | url=https://www.dsld.nlm.nih.gov/dsld/dailyvalue.jsp | access-date=16 May 2020 | archive-url=https://web.archive.org/web/20200407073956/https://dsld.nlm.nih.gov/dsld/dailyvalue.jsp | archive-date=7 April 2020 | url-status=dead }}</ref> A table of the old and new adult daily values is provided at [[Reference Daily Intake]].

==== Food sources ====
The main food sources for phosphorus are the same as those containing [[protein]], although proteins do not contain phosphorus. For example, milk, meat, and soya typically also have phosphorus. As a rule, if a diet has sufficient protein and calcium, the amount of phosphorus is probably sufficient.<ref>[[#Reference-Medline|"Phosphorus in diet"]]</ref>

==Precautions==
[[File:Phosphorus explosion.gif|thumb|Phosphorus explosion]]
Organic compounds of phosphorus form a broad class of materials; many are required for life, but some are highly toxic. Fluorophosphate [[ester]]s are among the most potent [[neurotoxin]]s known. A wide range of organophosphorus compounds are used for their toxicity as [[pesticides]] ([[herbicides]], [[insecticides]], [[fungicides]], etc.) and [[weapon]]ised as nerve agents against enemy humans. Most inorganic phosphates are relatively nontoxic and essential nutrients.<ref name="Greenwood" />

The white phosphorus allotrope presents a significant hazard because it ignites in the air and produces phosphoric acid residue. Chronic white phosphorus poisoning leads to necrosis of the jaw called "[[phossy jaw]]". White phosphorus is [[toxicity|toxic]], causing severe liver damage on ingestion and may cause a condition known as "Smoking Stool Syndrome".<ref>{{cite web| url = http://www.emedicine.com/EMERG/topic918.htm|title = CBRNE – Incendiary Agents, White Phosphorus (Smoking Stool Syndrome)| access-date = 2009-05-05}}</ref>

In the past, external exposure to elemental 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 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 name=":0">{{cite web |url=http://www.vnh.org/FM8285/Chapter/chapter9.html |title=US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries |access-date=2009-05-05 |archive-url=https://web.archive.org/web/20051122221207/http://www.vnh.org/FM8285/Chapter/chapter9.html |archive-date=November 22, 2005 |url-status=dead }}</ref>

The manual suggests instead "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#Ointment|ointments]] until it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns."{{#tag:ref|WP, (white phosphorus), exhibits chemoluminescence upon exposure to air and if there is any WP in the wound, covered by tissue or fluids such as blood serum, it will not glow until it is exposed to air, which requires a very dark room and dark-adapted eyes to see clearly|group=note}}<ref name=":0" /> 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.

In the workplace, people can be exposed to phosphorus by inhalation, ingestion, skin contact, and eye contact. The [[Occupational Safety and Health Administration]] (OSHA) has set the phosphorus exposure limit ([[Permissible exposure limit]]) in the workplace at 0.1&nbsp;mg/m<sup>3</sup> over an 8-hour workday. The [[National Institute for Occupational Safety and Health]] (NIOSH) has set a [[Recommended exposure limit]] (REL) of 0.1&nbsp;mg/m<sup>3</sup> over an 8-hour workday. At levels of 5&nbsp;mg/m<sup>3</sup>, phosphorus is [[IDLH|immediately dangerous to life and health]].<ref>{{Cite web|title = CDC - NIOSH Pocket Guide to Chemical Hazards - Phosphorus (yellow)|url = https://www.cdc.gov/niosh/npg/npgd0507.html|website = www.cdc.gov|access-date = 2015-11-21}}</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.|date = 1990|title = Methamphetamine 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, red and white phosphorus were designated by the United States [[Drug Enforcement Administration]] as [[DEA list of chemicals#List I chemicals|List I precursor chemicals]] under [[Code of Federal Regulations|21 CFR 1310.02]] effective on November 17, 2001.<ref name="66 CFR 52670">{{cite web| url = http://frwebgate.access.gpo.gov/cgi-bin/getdoc.cgi?dbname=2001_register&docid=01-26013-filed|title = 66 FR 52670—52675| date = 17 October 2001| access-date = 2009-05-05}}</ref> In the United States, handlers of red or white phosphorus are subject to stringent regulatory controls.<ref name="66 CFR 52670" /><ref name="21 CFR 1309">{{cite web| url =http://www.access.gpo.gov/nara/cfr/waisidx_06/21cfr1309_06.html| title =21 cfr 1309| access-date =2009-05-05| archive-url =https://web.archive.org/web/20090503063012/http://www.access.gpo.gov/nara/cfr/waisidx_06/21cfr1309_06.html| archive-date =2009-05-03}}</ref><ref name="CSA">{{cite web| url =http://www.usdoj.gov/dea/pubs/csa.html|title = 21 USC, Chapter 13 (Controlled Substances Act)| access-date = 2009-05-05}}</ref>

==See also==
*[[Hubbert peak theory]]

==Notes==
<references group=note/>

==Bibliography==
===References===
<references />

===General sources===
{{Refbegin}}<!--alphabetical, then by year-->
* {{Cite book |last=Beatty |first=Richard|title=Phosphorus|url=https://books.google.com/books?id=FHJIUJM1_JUC|publisher=Marshall Cavendish|date=2000|isbn=0-7614-0946-7}}
* {{cite news|newspaper=Miller-McCune|title=The Story of P(ee)|last=Burns|first=Melinda|date=10 February 2010|url=http://www.miller-mccune.com/science-environment/the-story-of-pee-8736/|access-date=2 February 2012|url-status=dead|archive-url=https://web.archive.org/web/20120107014214/http://www.miller-mccune.com/science-environment/the-story-of-pee-8736/|archive-date=7 January 2012}}
* {{wikicite|Corbridge, D.&nbsp;E.&nbsp;C. (1995) "Phosphorus: An Outline of its Chemistry, Biochemistry, and Technology" 5th Edition Elsevier: Amsterdam. {{ISBN|0-444-89307-5}}.|ref={{harvid|Corbridge|1995}}}}
* {{cite journal |last1=Cordell |first1=Dana |last2=Drangert |first2=Jan-Olof |last3=White |first3=Stuart |year=2009 |author-link1=Dana Cordell |title=The story of phosphorus: Global food security and food for thought |journal=Global Environmental Change |volume=19 |issue=2 |pages=292–305 |doi=10.1016/j.gloenvcha.2008.10.009 |bibcode=2009GEC....19..292C |s2cid=1450932 |issn=0959-3780}}
* {{Cite journal |last1=Cordell |first1=Dana |last2=White |first2=Stuart |date=2011-10-24 |title=Peak Phosphorus: Clarifying the Key Issues of a Vigorous Debate about Long-Term Phosphorus Security |journal=Sustainability |language=en |volume=3 |issue=10 |pages=2027–2049 |doi=10.3390/su3102027 |doi-access=free |issn=2071-1050|hdl=10453/31261 |hdl-access=free }}
* {{Ullmann|last1=Diskowski|first1=Herbert|last2=Hofmann|first2=Thomas|title=Phosphorus|doi=10.1002/14356007.a19_505}}
* {{cite book |last=Emsley |first=John |date=2000 |title=The Shocking history of Phosphorus: A biography of the Devil's Element. |location=London |publisher=MacMillan |isbn=0-333-76638-5 }}
* {{cite book |last=Emsley |first=John |date=7 January 2002 |title=The 13th Element: The Sordid Tale of Murder, Fire, and Phosphorus|url=https://books.google.com/books?id=D8IMOQAACAAJ|access-date=3 February 2012|publisher=John Wiley & Sons|isbn=978-0-471-44149-6}}
* {{wikicite<!--Cite preprint doesn't accept econ working papers, unfortunately-->|Heckenmüller, M.; Narita, D.; Klepper, G. (2014).  "[http://www.econstor.eu/bitstream/10419/90630/1/776834355.pdf Global availability of phosphorus and its implications for global food supply: An economic overview]".  Kiel Working Paper no.&nbsp;1897.  {{awrap|Accessed 11 May 2020.}}|ref={{harvid|Heckenmüller|Narita|Klepper|2014}}}}
* {{wikicite|[https://www.nlm.nih.gov/medlineplus/ency/article/002424.htm Phosphorus in diet: MedlinePlus Medical Encyclopedia]. Nlm.nih.gov (2011-11-07). Retrieved on 2011-11-19.|id=Medline}}
* {{cite book |last1=Parkes |first1=G. D. |last2=Mellor |first2=J. W. |date=1939 |title=Mellor's Modern Inorganic Chemistry |publisher=Longman's Green and Co. }}
* {{wikicite |Ruttenberg, K. C. [http://www.libraryindex.com/pages/3375/Phosphorus-Cycle.html Phosphorus Cycle – Terrestrial Phosphorus Cycle, Transport of Phosphorus], from ''Continents to the Ocean, The Marine Phosphorus Cycle''. {{webarchive|url=https://archive.today/20110713204340/http://www.libraryindex.com/pages/3375/Phosphorus-Cycle.html|date=July 13, 2011}}.|ref={{harvid|Ruttenberg}}}}
* {{cite book |last=Sommers |first=Michael&nbsp;A. |year=2007 |title=Phosphorus|publisher=Rosen Group|isbn=978-1-4042-1960-1|date=2007-08-15|url=https://archive.org/details/phosphorus0000somm/}}
* {{Ullmann |last1=Schrödter |first1=Klaus |first2=Gerhard|last2=Bettermann|first3=Thomas|last3=Staffel|first4=Friedrich|last4=Wahl|first5=Thomas|last5=Klein|first6=Thomas|last6=Hofmann|title=Phosphoric Acid and Phosphates|doi=10.1002/14356007.a19_465.pub3}}
* {{cite book |last=Threlfall |first=Richard E. |date=1951 |title=The Story of 100&nbsp;years of Phosphorus Making: 1851–1951 |location=Oldbury |publisher=Albright & Wilson Ltd }}
* {{cite book |last=Toy |first=Arthur D.&nbsp;F. |year=1975 |title=The Chemistry of Phosphorus|publisher=Pergamon|series=Texts in Inorganic Chemistry|volume=3|isbn=978-1-4831-4741-3|access-date=2013-10-22|url=https://archive.org/details/chemistryofphosp0003toya/}}
{{Refend}}

===Further reading===
* {{cite book |last1=Podger |first1=Hugh |date=2002 |title=Albright & Wilson. The Last 50&nbsp;years. |location=Studley |publisher=Brewin Books |isbn=1-85858-223-7}}
* [[Elizabeth Kolbert|Kolbert, Elizabeth]], "Elemental Need: Phosphorus helped save our way of life – and now threatens to end it", ''[[The New Yorker]]'', 6 March 2023, pp. 24–27. "[T]he world's phosphorus problem [arising from the element's exorbitant use in agriculture] resembles its carbon-dioxide problem, its plastics problem, its groundwater-use problem, its soil-erosion problem, and its nitrogen problem. The path humanity is on may lead to ruin, but, as of yet, no one has found a workable way back." (p. 27.)

{{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]]
[[Category:Chemical elements with body-centered cubic structure]]
[[Category:Peak resource production|Phosphorus]]