The isolation of white phosphorus in 1669 byHennig Brand marked the first "discovery" of an element not known since Antiquity. The name phosphorus is a reference to thegod of the Morning star inGreek mythology, inspired by the faint glow of white phosphorus when exposed tooxygen. This property is also at the origin of the termphosphorescence, meaning glow after illumination, although white phosphorus itself does not exhibit phosphorescence, butchemiluminescence caused by itsoxidation. Its hightoxicity makes exposure to white phosphorus very dangerous, while itsflammability andpyrophoricity can be weaponized in the form ofincendiaries. Red phosphorus is less dangerous and is used inmatches andfire retardants.
Most industrial production of phosphorus is focused on the mining and transformation of phosphate rock intophosphoric acid forphosphate-basedfertilisers. Phosphorus is an essential and oftenlimiting nutrient for plants, and while natural levels are normally maintained over time by thephosphorus cycle, it is too slow for the regeneration of soil that undergoesintensive cultivation. As a consequence, these fertilisers are vital to modern agriculture. The leading producers of phosphate ore in 2024 were China, Morocco, the United States and Russia, with two-thirds of the estimated exploitable phosphate reserves worldwide in Morocco alone. Other applications of phosphorus compounds includepesticides,food additives, anddetergents.
Phosphorus is essential to all known forms oflife, largely throughorganophosphates, organic compounds containing the phosphate ionPO3−4 as afunctional group. These includeDNA,RNA,ATP, andphospholipids, complex compounds fundamental to the functioning of allcells. The main component of bones and teeth,bone mineral, is a modified form ofhydroxyapatite, itself a phosphorus mineral.
Phosphorus was thefirst element to be "discovered", in the sense that it was not known since ancient times.[11] The discovery is credited to theHamburg alchemistHennig Brand in 1669, who was attempting to create the fabledphilosopher's stone.[12] To this end, he experimented withurine, which contains considerable quantities of dissolved phosphates from normal metabolism.[13] By letting the urine rot (a step later discovered to be unnecessary),[14] boiling it down to a paste, thendistilling it at a high temperature and leading the resulting vapours through water, he obtained a white, waxy substance that glowed in the dark and burned brilliantly. He named it inLatin:phosphorus mirabilis,lit. 'miraculous bearer of light'code: lat promoted to code: la. The word phosphorus itself (Ancient Greek:Φωσφόρος,romanized: Phōsphoros,lit. 'light-bearer') originates fromGreek mythology, where it references thegod of the morning star, also known as the planetVenus.[13][15]
Brand at first tried to keep the method secret,[16] but later sold the recipe for 200thalers toJohann Daniel Kraft [de] fromDresden.[13] Kraft toured much of Europe with it, includingLondon, where he met withRobert Boyle. The crucial fact that the substance was made from urine was eventually found out, andJohann Kunckel was able to reproduce it in Sweden in 1678. In 1680, Boyle also managed to make phosphorus and published the method of its manufacture.[13] He was the first to use phosphorus to ignitesulfur-tipped wooden splints, forerunners of modern matches,[17] and also improved the process by using sand in the reaction:
Finally, mixing the obtained calciummetaphosphate with groundcoal orcharcoal in an iron pot, and distilling phosphorus vapour out of aretort:
3 Ca(PO3)2 + 10 C → Ca3(PO4)2 + 10 CO + P4
This way, two-thirds of the phosphorus was turned into white phosphorus while one-third remained in the residue as calciumorthophosphate. Thecarbon monoxide produced during the reaction process was burnt off in aflare stack.
In 1609Inca Garcilaso de la Vega wrote the bookComentarios Reales in which he described many of the agricultural practices of the Incas prior to the arrival of the Spaniards and introduced the use ofguano as afertiliser. As Garcilaso described, the Incas near the coast harvested guano.[21] In the early 1800sAlexander von Humboldt introduced guano as a source of agricultural fertiliser to Europe after having discovered it in exploitable quantities on islands off the coast ofSouth America. It has been reported that, at the time of its discovery, the guano on some islands was over30 meters deep.[22] The guano had previously been used by theMoche people as a source of fertiliser by mining it and transporting it back toPeru by boat. International commerce in guano did not start until after 1840.[22] By the start of the 20th century guano had been nearly completely depleted and was eventually overtaken with the discovery of methods of production ofsuperphosphate.
Matches from 1828. The sulfur-tipped match is dipped into liquid containing white phosphorus, and ignites as it is pulled out of the bottle.
Early matches used white phosphorus in their composition, and were very dangerous due to both its toxicity and the way the match was ignited. The first striking match with a phosphorus head was invented byCharles Sauria in 1830. These matches (and subsequent modifications) were made with heads of white phosphorus, an oxygen-releasing compound (potassium chlorate,lead dioxide, or sometimesnitrate), and a binder. They were poisonous to the workers in manufacture, exposure to the vapours causing severenecrosis of the bones of the jaw, known as "phossy jaw".[23] Additionally, they were sensitive to storage conditions, toxic if ingested, and hazardous when accidentally ignited on a rough surface.[24][25] The very high risks for match workers was at the source of several notable early cases ofindustrial action, such as the 1888 LondonMatchgirls' strike.
The discovery of red phosphorus allowed for the development of matches that were both much safer to use and to manufacture, leading to the gradual replacement of white phoshphorus in matches. Additionally, around 1900 French chemists Henri Sévène and Emile David Cahen invented the modern strike-anywhere match, wherein the white phosphorus was replaced by phosphorus sesquisulfide (P4S3), a non-toxic and non-pyrophoric compound that ignites under friction. For a time these safer strike-anywhere matches were quite popular but in the long run they were superseded by the modern red phosphorus-based safety match. Following the implementation of these new manufacturing methods, production of white phosphorus matches was banned in several countries between 1872 and 1925,[26] and an internationaltreaty to this effect was signed following theBerne Convention (1906).[27]
Phosphate rock, which usually contains calcium phosphate, was first used in 1850 to make phosphorus. With the introduction of thesubmerged-arc furnace for phosphorus production byJames Burgess Readman in 1888[28] (patented 1889),[29] the use of bone-ash became obsolete.[30][31] 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.
White phosphorus shell explosion in France during the First World War (1918)
DuringWorld War II,Molotov cocktails made of phosphorus dissolved inpetrol were distributed in Britain to specially selected civilians as part of thepreparations for a potential invasion. The United States also developed the M15 white-phosphorus hand grenade, a precursor to theM34 grenade, while the British introduced the similarNo 77 grenade. These multipurpose grenades were mostly used for signaling and smoke screens, although they were also efficientanti-personnel weapons.[33] The difficulty of extinguishing burning phosphorus and the very severe burns it causes had a strong psychological impact on the enemy.[34] Phosphorusincendiary bombs were used on a large scale, notably todestroy Hamburg, the place where the "miraculous bearer of light" was first discovered.[15]
There are 22 knownisotopes of phosphorus,[35] ranging from26P to47P.[36] Only31P is stable and is therefore present at 100% abundance. The half-integernuclear spin and high abundance of31P makephosphorus-31 nuclear magnetic resonance spectroscopy a very useful analytical tool in studies of phosphorus-containing samples.
Tworadioactive isotopes of phosphorus have half-lives suitable for biological scientific experiments, and are used as radioactive tracers in biochemical laboratories.[37] These are:
33P, a beta-emitter (0.25 MeV) with a half-life of 25.4 days. It is used in life-science laboratories in applications in which lower energy beta emissions are advantageous such asDNA sequencing.
The high-energy beta particles from32P penetrate skin andcorneas and any32P ingested, inhaled, or absorbed is readily incorporated into bone andnucleic acids. For these reasons, personnel working with32P is required to wear lab coats, disposable gloves, and safety glasses, and avoid working directly over open containers.Monitoring personal, clothing, and surface contamination is also required. The high energy of the beta particles gives rise to secondary emission ofX-rays viaBremsstrahlung (braking radiation) in dense shielding materials such as lead. Therefore, the radiation must beshielded with low density materials such as water, acrylic or other plastic.[38]
Crystalline structures of the main phosphorus allotropes
White
Red
Violet
Black
Phosphorus has severalallotropes that exhibit very diverse properties.[39] The most useful and therefore common iswhite phosphorus, followed byred phosphorus. The two other main allotropes, violet and black phosphorus, have either a more fundamental interest or specialized applications.[40] Many other allotropes have been theorized and synthesized, with the search for new materials an active area of research.[41] Commonly mentioned "yellow phosphorus" is not an allotrope, but a result of the gradual degradation of white phosphorus into red phosphorus, accelerated by light and heat. This causes white phosphorus that is aged or otherwise impure (e.g. weapons-grade) to appear yellow.
White phosphorus is a soft, waxymolecular solid that is insoluble in water.[34] It is also very toxic, highlyflammable andpyrophoric, igniting in air at about 30 °C (303 K).[42] Structurally, it is composed ofP4tetrahedra. The nature of bonding in a givenP4 tetrahedron can be described byspherical aromaticity or cluster bonding, that is the electrons are highlydelocalized. This has been illustrated by calculations of the magnetically induced currents, which sum up to 29 nA/T, much more than in the archetypicalaromatic moleculebenzene (11 nA/T).[43] TheP4 molecule in the gas phase has a P-P bond length of 2.1994(3) Å as determined bygas electron diffraction.[43] White phosphorus exists in two crystalline forms named α (alpha) and β (beta), differing in terms of the relative orientation of the constituentP4 tetrahedra.[44][45] The α-form is most stable at room temperature and has acubic crystal structure. When cooled down to 195.2 K (−78.0 °C) it transforms into the β-form, turning into anhexagonal crystal structure. When heated up, the tetrahedral structure is conserved after melting at 317.3 K (44.2 °C) and boiling at 553.7 K (280.6 °C), before facingthermal decomposition at 1,100 K (830 °C) where it turns into gaseousdiphosphorus (P2).[46] This molecule contains a triple bond and is analogous toN2; it can also be generated as a transient intermediate in solution by thermolysis of organophosphorus precursor reagents.[47] At still higher temperatures,P2 dissociates into atomic P.[34]
White phosphorus exposed to air glows in the dark.
When exposed to air, white phosphorus faintly glows green and blue due tooxidation, a phenomenon best visible in the dark. This reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived moleculesHPO andP2O2 that both emit visible light.[48] However, in a pure-oxygen environment phosphorus does not glow at all, with the oxidation happening only in a range ofpartial pressures.[49] Derived from this phenomenon, the termsphosphors andphosphorescence have been loosely used to describe substances that shine in the dark. However, phosphorus itself is not phosphorescent butchemiluminescent, since it glows due to a chemical reaction and not the progressive reemission of previously absorbed light.[14]
Red phosphorus ispolymeric in structure. It can be viewed as a derivative ofP4 wherein one P-P bond is broken and one additional bond is formed with the neighbouring tetrahedron, resulting in chains ofP21 molecules linked byvan der Waals forces.[50] Red phosphorus may be formed by heating white phosphorus to 250 °C (523 K) in the absence of air or by exposing it to sunlight.[13] In this form phosphorus isamorphous, but can be crystallised upon further heating into violet phosphorus or fibrous red phosphorus depending on the reaction conditions. Red phosphorus is therefore not an allotrope in the strictest sense of the term, but rather an intermediate between other crystalline allotropes of phosphorus, and consequently most of its properties have a range of values. Freshly prepared, bright red phosphorus is highly reactive and ignites at about 300 °C (573 K).[51] After prolonged heating or storage, the color darkens; the resulting product is more stable and does not spontaneously ignite in air.[52]
Violet phosphorus or α-metallic phosphorus can be produced by day-long annealing of red phosphorus above 550 °C (823 K). In 1865,Johann Wilhelm Hittorf discovered that when phosphorus was recrystallised from moltenlead, a red/purple form is obtained. Therefore, this form is sometimes known as "Hittorf's phosphorus" .[53]
Black phosphorus or β-metallic phosphorus is the least reactive allotrope and the thermodynamically stable form below 550 °C (823 K). In appearance, properties, and structure, it resemblesgraphite, being black and flaky, a conductor of electricity, and having puckered sheets of linked atoms.[54][55][56] It is obtained by heating white phosphorus under high pressures (about 12,000 standard atmospheres or 1.2 gigapascals). It can also be produced at ambient conditions using metal salts, e.g. mercury, as catalysts.[57] Single-layer black phosphorus is calledphosphorene, and is therefore predictably analogous tographene.
In 2013, astronomers detected phosphorus inCassiopeia A, which confirmed that this element is produced insupernovae as a byproduct ofsupernova nucleosynthesis. The phosphorus-to-iron ratio in material from thesupernova remnant could be up to 100 times higher than in theMilky Way in general.[58] In 2020, astronomers analysedALMA andROSINA data from the massivestar-forming region AFGL 5142, to detect phosphorus-bearing molecules and how they could have been carried in comets to the early Earth.[59]
Phosphorus has a concentration in theEarth's crust of about one gram per kilogram (for comparison, copper is found at about 0.06 grams per kilogram). It is not found free in nature, but is widely distributed in manyminerals, usually as phosphates.[40] Inorganicphosphate rock, which is partially made ofapatite, is today the chief commercial source of this element.
The most prevalent compounds of phosphorus are derivatives of phosphate (PO3−4), a tetrahedral anion.[60] 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:
H3PO4 + H2O ⇌ H3O+ + H2PO−4 (Ka1 = 7.25×10−3)
H2PO−4 + H2O ⇌ H3O+ + HPO2−4 (Ka2 = 6.31×10−8)
HPO2−4 + H2O ⇌ H3O+ + PO3−4 (Ka3 = 3.98×10−13)
Food-gradephosphoric acid (additiveE338[61]) is used to acidify foods and beverages such as variouscolas and jams, providing a tangy or sour taste.[62] The phosphoric acid also serves as apreservative.[63] Soft drinks containing phosphoric acid, includingCoca-Cola, are sometimes calledphosphate sodas or phosphates. Phosphoric acid in soft drinks has the potential to cause dental erosion,[64] as well as contribute to the formation ofkidney stones, especially in those who have had kidney stones previously.[65]
With metalcations, 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 (HPO2−4).
Calcium phosphates in particular are widespread compounds with many applications. Among them, they are used to improve the characteristics of processed meat andcheese, inbaking powder, and in toothpaste.[62]. Two of the most relevant among them aremonocalcium phosphate, anddicalcium phosphate.
Phosphate exhibits a tendency to form chains and rings containing P-O-P bonds. Many polyphosphates are known, includingATP. Polyphosphates arise by dehydration of hydrogen phosphates such asHPO2−4 andH2PO−4. For example, the industrially important pentasodium triphosphate (also known assodium tripolyphosphate, STPP) is produced industrially by the megatonne by thiscondensation reaction:
2 Na2HPO4 + NaH2PO4 → Na5P3O10 + 2 H2O
Sodium triphosphate is used in laundry detergents in some countries, but banned for this use in others.[52] This compoundsoftens the water to enhance the performance of the detergents and to prevent pipe and boiler tubecorrosion.[66]
Phosphorusoxoacids 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.[34] Although many oxoacids of phosphorus are formed, only nine are commercially important. Among them, hypophosphorous, phosphorous and orthophosphoric acid are particularly important.
Phosphorus pentoxide (P4O10) is theacid anhydride of phosphoric acid, but several intermediates between the two are known. This waxy white solid reacts vigorously with water. Similarly,phosphorus trioxide (P4O6, also called tetraphosphorus hexoxide) is the anhydride ofP(OH)3, the minor tautomer of phosphorous acid. The structure ofP4O6 is like that ofP4O10 without the terminal oxide groups. Mixed oxyhalides and oxyhydrides of phosphorus(III) are almost unknown. Meanwhile, phosphorus forms a wide range of sulfides, where the phosphorus can be in P(V), P(III) or other oxidation states. However, only two of them commercially significant.Phosphorus pentasulfide (P4S10) has a structure analogous toP4O10, and is used in the manufacture of additives and pesticides.[67] The three-fold symmetricPhosphorus sesquisulfide (P4S3) is used instrike-anywhere matches.
Phosphorushalides can have as oxidation state +3 in the case of trihalides and +5 for pentahalides andchalcoalides, but also +2 for disphosphorus tetrahalides. All four symmetrical trihalides are well known: gaseousPF3, the yellowish liquidsPCl3 andPBr3, and the solidPI3. These materials are moisture sensitive, hydrolysing to givephosphorous acid. The trichloride, a common reagent used for the manufacture of pesticides, is produced by chlorination of white phosphorus. The trifluoride is produced from the trichloride by halide exchange.PF3 is toxic because it binds tohaemoglobin.
Most phosphorus pentahalides are common compounds.PF5 is a colourless gas and the molecules have atrigonal bipyramidal geometry. With fluoride, it formsPF−6, ananion that isisoelectronic withSF6.PCl5 is a colourless solid which has an ionic formulation ofPCl+4PCl−6, but adopts a trigonal bipyramidal geometry when molten or in the vapour phase.[34] Both the pentafluoride and the pentachloride areLewis acids. Meanwhile,PBr5 is an unstable solid formulated asPBr+4Br−.PI5 is not known.[34]
The most important phosphorusoxyhalide isphosphorus oxychloride (POCl3), which is approximately tetrahedral. It is prepared fromPCl3 and used in the manufacture of plasticizers. Phosphorus can also form thiohalides such asPSCl3, and in rare cases selenohalides.
The PN moleculephosphorus mononitride is considered unstable, but is a product of crystallinetriphosphorus pentanitride decomposition at 1,100 K (830 °C). Similarly,H2PN is considered unstable, and phosphorus nitride halogens likeF2PN,Cl2PN,Br2PN, andI2PN oligomerise into cyclicpolyphosphazenes. For example, compounds of the formula(PNCl2)n exist mainly as rings such as thetrimerhexachlorophosphazene. The phosphazenes arise by treatment of phosphorus pentachloride with ammonium chloride:
PCl5 + NH4Cl → 1/n (NPCl2)n + 4 HCl
When the chloride groups are replaced byalkoxide (RO−), a family of polymers is produced with potentially useful properties.[68]
A wide variety of compounds which contain the containing the phosphide ionP3− exist, both withmain-group elements and withmetals. They often exhibit complex structures, where phosphorus has the -3 oxidation state. Metal phosphides arise by reaction of metals with red phosphorus. Thealkali metals (group 1) andalkaline earth metals (group 2) can also form compounds such asNa3P7. These compounds react with water to formphosphine.[34] Some phosphide minerals are also known, like(Fe,Ni)2P and(Fe,Ni)3P, but they are very rare on Earth, most instances occurring iniron-nickel meteorites.
Phosphine (PH3) and its organic derivatives are structural analogues ofammonia (NH3), but the bond angles at phosphorus are closer to 90° for phosphine and its organic derivatives. It is an ill-smelling and toxic gas, produced by hydrolysis ofcalcium phosphide (Ca3P2). Unlike ammonia, phosphine is oxidised by air. Phosphine is also far less basic than ammonia. Other phosphines are known which contain chains of up to nine phosphorus atoms and have the formulaPnHn+2.[34] The highly flammable gasdiphosphine (P2H4) is an analogue ofhydrazine.
Compounds with P-C and P-O-C bonds are often classified as organophosphorus compounds. They are widely used commercially. TheP3+ serves as a source ofPCl3 in routes to organophosphorus(III) compounds. For example, it is the precursor totriphenylphosphine:
PCl3 + 6 Na + 3 C6H5Cl → P(C6H5)3 + 6 NaCl
Treatment of phosphorus trihalides with alcohols andphenols gives phosphites, e.g.triphenylphosphite:
While many organic compounds of phosphorus are required for life, some are highly toxic. A wide range of organophosphorus compounds are used for their toxicity aspesticides andweaponised asnerve agents.[34] Some notable examples includesarin,VX orTabun. Fluorophosphateesters (like sarin) are among the most potentneurotoxins known.
Symmetric phosphorus(III) trithioesters (e.g.P(SMe)3) can be produced from the reaction ofwhite phosphorus and the correspondingdisulfide, or phosphorus(III) halides andthiolates. Unlike the corresponding esters, they do not undergo a variant of theMichaelis-Arbuzov reaction with electrophiles, instead reverting to another phosphorus(III) compound through asulfonium intermediate.[69]
These compounds generally feature P–P bonds.[34] Examples include catenated derivatives of phosphine and organophosphines. Compounds containing P=P double bonds have also been observed, although they are rare.
Inorganic phosphorus in the form of the phosphatePO3−4 is required for all known forms oflife.[70] Phosphorus plays a major role in the structural framework ofDNA andRNA. Living cells use phosphate to transport cellular energy withadenosine triphosphate (ATP), necessary for every cellular process that uses energy. ATP is also important forphosphorylation, a key regulatory event in cells. Every living cell is encased in a membrane that separates it from its surroundings. Cellular membranes are composed of a phospholipid matrix and proteins, typically in the form of a bilayer.Phospholipids are derived fromglycerol with two of the glycerol hydroxyl (OH) protons replaced by fatty acids as anester, and the third hydroxyl proton has been replaced with phosphate bonded to another alcohol.[71]
The main component of bone ishydroxyapatite 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 materialfluorapatite:[34]
Ca5(PO4)3OH + F− → Ca5(PO4)3F + OH−
An average adult human contains about 0.7 kilograms (1.5 lb) of phosphorus, about 85–90% in bones and teeth in the form ofapatite, and the remainder in soft tissues and extracellular fluids. The phosphorus content increases from about 0.5% by mass in infancy to 0.65–1.1% by mass in adults. In comparison, average phosphorus concentration in the blood is about 0.4 g/L; about 70% of that is organic and 30% inorganic phosphates.[72]
The main food sources for phosphorus are the same as those containingprotein, although proteins themselves do not contain phosphorus. For example, milk, meat, and soya typically also have phosphorus. Generally, if a diet includes sufficient protein and calcium, the amount of phosphorus is sufficient.[73]
According to theU.S. Institute of Medicine, the estimated average requirement for phosphorus for people ages 19 and up is 580 mg/day. The RDA is 700 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 mg/day. For people ages 1–18 years, the RDA increases with age from 460 to 1250 mg/day. As for safety, the IOM setstolerable upper intake level for phosphorus at 4000 mg/day. Collectively, these values are referred to as theDietary Reference Intake.[74] TheEuropean 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.[75] AI and UL are defined the same as in the United States. For people ages 15 and older, including pregnancy andlactation, the AI is set at 550 mg/day. For children ages 4–10, the AI is 440 mg/day, and for ages 11–17 it is 640 mg/day. These AIs are lower than the U.S. RDAs. In both systems, teenagers need more than adults.[76] The EFSA reviewed the same safety question and decided that there was not sufficient information to set a UL.[77]
Phosphorus deficiency may be caused bymalnutrition, by failure to absorb phosphate, and by metabolic syndromes that draw phosphate from the blood (such as inrefeeding syndrome after malnutrition[78]) or passing too much of it into the urine. All are characterised byhypophosphatemia, 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 ofATP. 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.[79]
Phosphorus is an essential plant nutrient (the most often limiting nutrient, afternitrogen),[80] and the bulk of all phosphorus production is in concentrated phosphoric acids foragriculturefertilisers, containing as much as 70% to 75% P2O5. That led to large increase inphosphate (PO43−) production in the second half of the 20th century.[81] 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 ofplant roots, formation of seeds and flowers, and other important factors effecting overall plant health and genetics.[80] Heavy use of phosphorus fertilizers and their runoff have resulted ineutrophication (overenrichment) ofaquatic ecosystems.[82][83]
Natural phosphorus-bearing compounds are mostly inaccessible to plants because of the low solubility and mobility in soil.[84] 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 naturalphosphorus 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.[85] Fertiliser is often in the form of superphosphate of lime, a mixture of calcium dihydrogen phosphate (Ca(H2PO4)2), and calcium sulfate dihydrate (CaSO4·2H2O) produced reacting sulfuric acid and water withcalcium phosphate.
Processing phosphate minerals with sulfuric acid for obtaining fertiliser is so important to the global economy that this is the primary industrial market forsulfuric acid and the greatest industrial use of elementalsulfur.[86]
Means of commercial phosphorus production besides mining are few because thephosphorus cycle does not include significant gas-phase transport.[87] The predominant source of phosphorus in modern times is phosphate rock (as opposed to the guano that preceded it).
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 ofChina. In 2019, the US produced 10 percent of the world's phosphate rock.[88]
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 givephosphoric acid. Phosphoric acid is then neutralized to give various phosphate salts, which comprise fertilizers. In the wet process, phosphorus does not undergo redox.[89] About five tons ofphosphogypsum waste are generated per ton of phosphoric acid production. Annually, the estimated generation of phosphogypsum worldwide is 100 to 280 Mt.[90]
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 asubmerged-arc furnace which is energy intensive.[89] Presently, about 1,000,000short tons (910,000 t) of elemental phosphorus is produced annually.Calcium phosphate (asphosphate rock), mostly mined in Florida and North Africa, can be heated to 1,200–1,500 °C with sand, which is mostlySiO2, andcoke to produceP4. TheP4 product, being volatile, is readily isolated:[91]
4 Ca5(PO4)3F + 18 SiO2 + 30 C → 3 P4 + 30 CO + 18 CaSiO3 + 2 CaF2
2 Ca3(PO4)2 + 6 SiO2 + 10 C → 6 CaSiO3 + 10 CO + P4
Side products from the thermal process includeferrophosphorus, a crude form ofFe2P, resulting from iron impurities in the mineral precursors. The silicateslag is a useful construction material. The fluoride is sometimes recovered for use inwater 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.[92]
Annual global phosphate rock production (megatonnes per yr), 1994–2022 (data from US Geological Survey)[93]
Phosphorus comprises about 0.1% by mass of theEarth's crust.[94][95][96] However, only concentrated forms collectively referred to asphosphate rock or phosphorite are exploitable, and are not evenly distributed across the Earth.[34] Unprocessed phosphate rock has a concentration of 1.7–8.7% phosphorus by mass (4–20% phosphorus pentoxide). The world's total commercial phosphate reserves and resources are estimated in amounts of phosphate rock, which in practice includes over 300 ores of different origin, composition, and phosphate content. "Reserves" refers to the amount assumed recoverable at current market prices and "resources" refers to estimated amounts of such a grade or quality that they have reasonable prospects for economic extraction.[97][98] Mining is currently the only cost-effective method for the production of phosphorus. Hence, a shortage in rock phosphate or significant price increases might negatively affect the world'sfood security.[99]
Global distribution of commercial reserves of rock phosphate in 2016[100]
The countries estimated to have the biggest phosphate rock commercial reserves (in billion metric tons) areMorocco (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), andFinland (1.0).[101][102][103] 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.[104] According to some researchers, Earth's commercial and affordable phosphorus reserves are expected to be depleted in 50–100 years.[105]
In 2023, theUnited 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.[102] Assuming zero growth, the reserves would thus last for around 300 years. This broadly confirms a 2010International Fertilizer Development Center (IFDC) report that global reserves would last for several hundred years.[106][107] Phosphorus reserve figures are intensely debated.[97][108][109] Gilbert suggest that there has been little external verification of the estimate.[110] A 2014 review[111] 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".
A phosphate removal sewage treatment station inYorkshire, England
Reducing agricultural runoff and soil erosion can slow the frequency with which farmers have to reapply phosphorus to their fields. Agricultural methods such asno-till farming,terracing,contour tilling, and the use ofwindbreaks have been shown to reduce the rate of phosphorus depletion from farmland, though do not completely remove the need for periodic fertilizer application. Strips of grassland or forest between arable land and rivers can also greatly reduce losses of phosphate and other nutrients.[112]
Sewage treatment plants that have adedicated phosphorus removal step produce phosphate-richsewage sludge that can then betreated to extract phosphorus from it. This is done byincinerating the sludge and recovering the resulting ash.[113] Another approach lies into the recovery of phosphorus-rich materials such asstruvite from waste processing plants, which is done by adding magnesium to the waste.[110] However, the technologies currently in use are not yet cost-effective, given the current price of phosphorus on the world market.[114][115]
Phosphorus is also an important component insteel production, in the making ofphosphor bronze, and in many other related products.[118][119] 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 higherhydrogen embrittlement resistance than normal copper.[120]Phosphate conversion coating is a chemical treatment applied to steel parts to improve their corrosion resistance.
Match striking surface made of a mixture of red phosphorus, glue and ground glass. The glass powder is used to increase the friction.
Safety matches are very difficult to ignite on any surface other than a special striker strip. The strip contains non-toxic red phosphorus and the match headpotassium chlorate, an oxygen-releasing compound. When struck, small amounts ofabrasion from match head and striker strip are mixed intimately to make a small quantity ofArmstrong'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.[51][121]
Phosphorus is adopant inN-type semiconductors used in high-power electronics andsemiconductor detectors.[126] In this context, phosphorus is not present at the start of the process, but rather created directly out of silicon during the manufacture of the devices. This is done by neutrontransmutation doping, a method based on the conversion of the30Si into31P byneutron capture andbeta decay as follows:
In practice, the silicon is typically placed near or inside anuclear reactor generating neutrons. As neutrons pass through the silicon, phosphorus atoms are produced by transmutation. This doping method is far less common than diffusion or ion implantation, but it has the advantage of creating an extremely uniform dopant distribution.[127][128]
Elemental phosphorus poses by far the greatest danger in its white form, red phosphorus being relatively nontoxic.[129] In the past, external exposure to white phosphorus was treated by washing the affected area with 2%copper(II) sulfate solution to form harmless compounds that are then washed away. According to 2009United States Navy guidelines:[130]
Cupric (copper) sulfate has been used by U.S. personnel in the past and is still being used by some nations. However, copper sulfate is toxic and its use will be discontinued. Copper sulfate may produce kidney and cerebral toxicity as well asintravascular hemolysis.
Instead, the manual suggests:
[...] 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. Promptlydebride 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-basedointments until it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns.
Because of its common use as arodenticide, there are documented medical reports of white phosphorus ingestion and its effects, especially on children.[131] These cases can present very characteristic symptoms, such as garlic-smelling, smoking and luminescent vomit and stool, the latter sometimes called "Smoking Stool Syndrome". It is absorbed by both the gastrointestinal tract and the respiratory mucosa, to whose it causes serious damage. The acute lethal dose has been estimated at around 1 mg/kg, the very small amount resulting in many cases proving fatal, either because of rapid cardiovascular arrest or through the following systemic toxicity.[131]
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