Phosphorus

This article is about the chemical element. For other uses, seePhosphorus (disambiguation).

Phosphorusis achemical element;it hassymbolPandatomic number15. Elemental phosphorus exists in two major forms,white phosphorusandred phosphorus,but because it is highlyreactive,phosphorus is never found as a free element on Earth. It has a concentration in the Earth's crust of about one gram per kilogram (compare copper at about 0.06 grams). In minerals, phosphorus generally occurs asphosphate.

Phosphorus,₁₅P

Forms of phosphorus Waxy white Light red Dark red and violet Black
Phosphorus
Pronunciation	/ˈfɒsfərəs/​(FOS-fər-əs)
Allotropes	white, red, violet, black and others (seeAllotropes of phosphorus)
Appearance	white, red and violet are waxy, black is metallic-looking
Standard atomic weightAr°(P)
	30.973761998±0.000000005 30.974±0.001(abridged)
Abundance
in theEarth's crust	5.2 (silicon = 100)
Phosphorus in theperiodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson N ↑ P ↓ As silicon←phosphorus→sulfur
Atomic number(Z)	15
Group	group 15 (pnictogens)
Period	period 3
Block	p-block
Electron configuration	[Ne] 3s23p3
Electrons per shell	2, 8, 5
Physical properties
PhaseatSTP	solid
Melting point	white: 317.3K​(44.15 °C, ​111.5 °F) red: ∼860 K (∼590 °C, ∼1090 °F)
Boiling point	white: 553.7 K ​(280.5 °C, ​536.9 °F)
Sublimation point	red: ≈689.2–863 K ​(≈416–590 °C, ​≈780.8–1094 °F) violet: 893 K (620 °C, 1148 °F)
Density(nearr.t.)	white: 1.823 g/cm3 red: ≈2.2–2.34 g/cm3 violet: 2.36 g/cm3 black: 2.69 g/cm3
Heat of fusion	white: 0.66kJ/mol
Heat of vaporisation	white: 51.9 kJ/mol
Molar heat capacity	white: 23.824 J/(mol·K)
Vapour pressure(white) P(Pa) 1 10 100 1 k 10 k 100 k atT(K) 279 307 342 388 453 549
Vapour pressure(red) P(Pa) 1 10 100 1 k 10 k 100 k atT(K) 455 489 529 576 635 704
Atomic properties
Oxidation states	−3,−2, −1, 0,+1,+2,+3,+4,+5(a mildlyacidicoxide)
Electronegativity	Pauling scale: 2.19
Ionisation energies	1st: 1011.8 kJ/mol 2nd: 1907 kJ/mol 3rd: 2914.1 kJ/mol (more)
Covalent radius	107±3pm
Van der Waals radius	180 pm
Spectral linesof phosphorus
Other properties
Natural occurrence	primordial
Crystal structure	α-white: ​body-centred cubic(bcc) (cI232)
Lattice constant	a= 1.869 nm (at 20 °C)
Crystal structure	black: ​orthorhombic(oS8)
Lattice constants	a= 0.33137 nm b= 1.0477 nm c= 0.43755 nm (at 20 °C)
Thermal conductivity	white: 0.236 W/(m⋅K) black: 12.1 W/(m⋅K)
Magnetic ordering	white, red, violet, black:diamagnetic
Molar magnetic susceptibility	−20.8×10−6cm3/mol (293 K)
Bulk modulus	white: 5 GPa red: 11 GPa
CAS Number	7723-14-0 (red) 12185-10-3 (white)
History
Discovery	Hennig Brand(1669)
Recognised as an element by	Antoine Lavoisier(1777)
Isotopes of phosphorus
Main isotopes Decay abun­dance half-life(t1/2) mode pro­duct 31P 100% stable 32P trace 14.269 d β− 32S 33P trace 25.35 d β− 33S

Elemental phosphorus was first isolated as white phosphorus in 1669. In white phosphorus, phosphorus atoms are arranged in groups of 4, written as P₄.White phosphorus emits a faint glow when exposed tooxygen—hence a name, taken from Greek mythology,Φωσφόροςmeaning 'light-bearer' (LatinLucifer), referring to the "Morning Star",the planetVenus.The termphosphorescence,meaning glow after illumination, has its origin in phosphorus, although phosphorus itself does not exhibit phosphorescence: phosphorus glows due tooxidationof the white (but not red) phosphorus—a process now calledchemiluminescence.Phosphorus is classified as apnictogen,together withnitrogen,arsenic,antimony,bismuth,andmoscovium.

Phosphorus is an element essential to sustaininglifelargely throughphosphates,compounds containing the phosphate ion, PO₄³⁻.Phosphates are a component ofDNA,RNA,ATP,andphospholipids,complex compounds fundamental tocells.Elemental phosphorus was first isolated from humanurine,andbone ashwas 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 asfertilisers.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 includeorganophosphorus compoundsindetergents,pesticides,andnerve agents.

Characteristics

Allotropes

Main article:Allotropes of phosphorus

Phosphorus has severalallotropesthat exhibit strikingly diverse properties.The two most common allotropes are white phosphorus and red phosphorus.

For both pure and applied uses, the most important allotrope iswhite phosphorus,often abbreviated WP. White phosphorus is a soft, waxymolecular solidcomposed ofP
₄tetrahedra.ThisP
₄tetrahedron is also present in liquid and gaseous phosphorus up to the temperature of 800 °C (1,500 °F; 1,100 K) when it starts decomposing toP
₂molecules.The nature of bonding in thisP
₄tetrahedron can be described byspherical aromaticityor 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 archetypicalaromaticmoleculebenzene(11 nA/T).

Crystalline structures of some phosphorus allotropes

White

Red

Violet

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 195.2 K (−78.0 °C), it transforms into β-form, which has hexagonal crystal structure. These forms differ in terms of the relative orientations of the constituent P₄tetrahedra.

White phosphorus is the least stable, the most reactive, the mostvolatile,the leastdenseand 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 highlyflammableandpyrophoric(self-igniting) in air; it faintly glows green and blue in the dark when exposed to oxygen. The autoxidation commonly coats samples with whitephosphorus pentoxide(P
₄O
₁₀): P₄tetrahedra, but with oxygen inserted between the phosphorus atoms and at the vertices. White phosphorus is anapalmadditive, and the characteristic odour of combustion is garlicky. White phosphorus is insoluble in water but soluble in carbon disulfide.

Thermal decompositionof P₄at 1100 K givesdiphosphorus,P₂.This species is not stable as a solid or liquid. The dimeric unit contains a triple bond and is analogous to N₂.It can also be generated as a transient intermediate in solution by thermolysis of organophosphorus precursor reagents.At still higher temperatures, P₂dissociates into atomic P.

Properties of some allotropes of phosphorus

Form	white(α)	white(β)	red	violet	black
Symmetry	Body-centred cubic	Triclinic	Amorphous	Monoclinic	Orthorhombic
Pearson symbol		aP24		mP84	oS8
Space group	I43m	P1No.2		P2/c No.13	Cmce No.64
Density(g/cm3)	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 phosphorusis polymeric in structure. It can be viewed as a derivative of P₄wherein one P-P bond is broken, and one additional bond is formed with the neighbouring tetrahedron resulting in chains of P₂₁molecules linked byvan der Waals forces.Red phosphorus may be formed by heating white phosphorus to 250 °C (482 °F) or by exposing white phosphorus to sunlight.Phosphorus after this treatment isamorphous.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 300 °C (572 °F),though it is more stable than white phosphorus, which ignites at about 30 °C (86 °F).After prolonged heating or storage, the color darkens (see infobox images); the resulting product is more stable and does not spontaneously ignite in air.

Violet phosphorusis a form of phosphorus that can be produced by day-long annealing of red phosphorus above 550 °C. In 1865,Hittorfdiscovered that when phosphorus was recrystallised from moltenlead,a red/purple form is obtained. Therefore, this form is sometimes known as "Hittorf's phosphorus" (or violet or α-metallic phosphorus).

Black phosphorusis the least reactive allotrope and the thermodynamically stable form below 550 °C (1,022 °F). It is also known as β-metallic phosphorus and has a structure somewhat resembling that ofgraphite.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.In appearance, properties, and structure, it resemblesgraphite,being black and flaky, a conductor of electricity, and has puckered sheets of linked atoms.

Another form, scarlet phosphorus, is obtained by allowing a solution of white phosphorus incarbon disulfideto evaporate insunlight.

Chemiluminescence

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 Boylein 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;there is only a range ofpartial pressuresat which it does. Heat can be applied to drive the reaction at higher pressures.

In 1974, the glow was explained by R. J. van Zee and A. U. Khan.A reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived molecules HPO andP
₂O
₂that 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,phosphorsandphosphorescencewere used loosely to describe substances that shine in the dark without burning. Although the termphosphorescenceis derived from phosphorus, the reaction that gives phosphorus its glow is properly calledchemiluminescence(glowing due to a cold chemical reaction), not phosphorescence (re-emitting light that previously fell onto a substance and excited it).

Isotopes

Main article:Isotopes of phosphorus

There are 22 knownisotopesof phosphorus,ranging from²⁶
Pto⁴⁷
P.Only³¹
Pis stable and is therefore present at 100% abundance. The half-integernuclear spinand high abundance of³¹P makephosphorus-31 NMRspectroscopy a very useful analytical tool in studies of phosphorus-containing samples.

Tworadioactive isotopesof phosphorus have half-lives suitable for biological scientific experiments. These are:

• ³²
  P,abeta-emitter (1.71 MeV) with ahalf-lifeof 14.3 days, which is used routinely in life-science laboratories, primarily to produceradiolabeledDNA and RNAprobes,e.g. for use inNorthern blotsorSouthern blots.
• ³³
  P,a beta-emitter (0.25 MeV) with a half-life of 25.4 days. It is used in life-science laboratories in applications in which lower energy beta emissions are advantageous such asDNAsequencing.

The high-energy beta particles from³²
Ppenetrate skin andcorneasand any³²
Pingested, inhaled, or absorbed is readily incorporated into bone andnucleic acids.For these reasons,Occupational Safety and Health Administrationin the United States, and similar institutions in other developed countries require personnel working with³²
Pto wear lab coats, disposable gloves, and safety glasses or goggles to protect the eyes, and avoid working directly over open containers.Monitoringpersonal, clothing, and surface contamination is also required.Shieldingrequires special consideration. The high energy of the beta particles gives rise to secondary emission ofX-raysviaBremsstrahlung(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.

Occurrence

See also:Category:Phosphate minerals

Universe

In 2013, astronomers detected phosphorus inCassiopeia A,which confirmed that this element is produced insupernovaeas a byproduct ofsupernova nucleosynthesis.The phosphorus-to-ironratio in material from thesupernova remnantcould be up to 100 times higher than in theMilky Wayin general.

In 2020, astronomers analysedALMAandROSINAdata from the massivestar-forming regionAFGL 5142, to detect phosphorus-bearing molecules and how they are carried in comets to the early Earth.

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 manyminerals,usually as phosphates.Inorganicphosphate rock,which is partially made ofapatite(a group of minerals being, generally, pentacalcium triorthophosphate fluoride (hydroxide)), is today the chief commercial source of this element. According to theUS Geological Survey (USGS),about 50 percent of the global phosphorus reserves are inAmazighnations likeMorocco,AlgeriaandTunisia.85% of Earth's known reserves are inMoroccowith smaller deposits inChina,Russia,Florida,Idaho,Tennessee,Utah,and elsewhere.Albright and Wilsonin the UK and theirNiagara Fallsplant, for instance, were using phosphate rock in the 1890s and 1900s from Tennessee, Florida, and theÎles du Connétable(guanoisland sources of phosphate); by 1950, they were using phosphate rock mainly from Tennessee and North Africa.

Organic sources, namelyurine,bone ashand (in the latter 19th century)guano,were historically of importance but had only limited commercial success.As urine contains phosphorus, it has fertilising qualities which are still harnessed today in some countries, includingSweden,using methods forreuse 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 ofsewageorsewage sludge.

Compounds

See also:Phosphorus halides,Category:Phosphate minerals,andCategory:Phosphorus compounds

Phosphorus(V)

The tetrahedral structure of P₄O₁₀and P₄S₁₀

The most prevalent compounds of phosphorus are derivatives of phosphate (PO₄³⁻), a tetrahedral anion.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⁺+ H₂PO₄⁻K_a1 = 7.25×10⁻³

H₂PO₄⁻+ H₂O ⇌ H₃O⁺+ HPO₄²⁻K_a2= 6.31×10⁻⁸

HPO₄²⁻+ H₂O ⇌ H₃O⁺+ PO₄³⁻K_a3= 3.98×10⁻¹³

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 as HPO₄²⁻and H₂PO₄⁻.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

Phosphorus pentoxide(P₄O₁₀) is theacid anhydrideof phosphoric acid, but several intermediates between the two are known. This waxy white solid reacts vigorously with water.

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 (HPO₄²⁻).

PCl₅andPF₅are common compounds. PF₅is a colourless gas and the molecules havetrigonal bipyramidalgeometry. PCl₅is a colourless solid which has an ionic formulation of PCl₄⁺PCl₆⁻,but adopts thetrigonal bipyramidalgeometry when molten or in the vapour phase.PBr₅is an unstable solid formulated as PBr₄⁺Br⁻andPI₅is not known.The pentachloride and pentafluoride areLewis acids.With fluoride, PF₅forms PF₆⁻,ananionthat isisoelectronicwith SF₆.The most important oxyhalide isphosphorus oxychloride,(POCl₃), which is approximately tetrahedral.

Before extensive computer calculations were feasible, it was thought that bonding in phosphorus(V) compounds involveddorbitals. Computer modeling ofmolecular orbital theoryindicates that this bonding involves only s- and p-orbitals.

Phosphorus(III)

All four symmetrical trihalides are well known: gaseousPF₃,the yellowish liquidsPCl₃andPBr₃,and the solidPI₃.These materials are moisture sensitive, hydrolysing to givephosphorous acid.The trichloride, a common reagent, is produced by chlorination of white phosphorus:

P₄+ 6 Cl₂→ 4 PCl₃

The trifluoride is produced from the trichloride by halide exchange. PF₃is toxic because it binds tohaemoglobin.

Phosphorus(III) oxide,P₄O₆(also called tetraphosphorus hexoxide) is the anhydride of P(OH)₃,the minor tautomer of phosphorous acid. The structure of P₄O₆is like that of P₄O₁₀without the terminal oxide groups.

Phosphorus(I) and phosphorus(II)

A stablediphosphene,a derivative of phosphorus(I)

These compounds generally feature P–P bonds.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

Phosphidesarise by reaction of metals with red phosphorus. The alkali metals (group 1) and alkaline earth metals can form ionic compounds containing thephosphideion, P³⁻.These compounds react with water to formphosphine.Otherphosphides,for example Na₃P₇,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.Schreibersiteis a naturally occurring metal-rich phosphide found in meteorites. The structures of the metal-rich and phosphorus-rich phosphides can be complex.

Phosphine(PH₃) and its organic derivatives (PR₃) are structural analogues of ammonia (NH₃), 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 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 formula P_nH_n+2.The highly flammable gasdiphosphine(P₂H₄) is an analogue ofhydrazine.

Oxoacids

Phosphorusoxoacidsare 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.Although many oxoacids of phosphorus are formed, only nine are commercially important, and three of them,hypophosphorous acid,phosphorous acid,andphosphoric acid,are particularly important.

Oxidation state	Formula	Name	Acidic protons	Compounds
+1	HH2PO2	hypophosphorous acid	1	acid, salts
+3	H3PO3	phosphorous acid (phosphonic acid)	2	acid, salts
+3	HPO2	metaphosphorous acid	1	salts
+4	H4P2O6	hypophosphoric acid	4	acid, salts
+5	(HPO3)n	metaphosphoric acids	n	salts (n= 3,4,6)
+5	H(HPO3)nOH	polyphosphoric acids	n+2	acids, salts (n= 1-6)
+5	H5P3O10	tripolyphosphoric acid	3	salts
+5	H4P2O7	pyrophosphoric acid	4	acid, salts
+5	H3PO4	(ortho)phosphoric acid	3	acid, salts

Nitrides

The PN molecule is considered unstable, but is a product of crystallinephosphorus nitridedecomposition at 1100 K. Similarly, H₂PN is considered unstable, and phosphorus nitride halogens like F₂PN, Cl₂PN, Br₂PN, and I₂PN oligomerise into cyclicpolyphosphazenes.For example, compounds of the formula (PNCl₂)_nexist 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.

Sulfides

Main article: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 P₄S₃is used in strike-anywhere matches. P₄S₁₀and P₄O₁₀have analogous structures.Mixed oxyhalides and oxyhydrides of phosphorus(III) are almost unknown.

Organophosphorus compounds

Main article:organophosphorus compounds

Compounds with P-C and P-O-C bonds are often classified as organophosphorus compounds. They are widely used commercially. The PCl₃serves as a source of P³⁺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 andphenolsgives phosphites, e.g.triphenylphosphite:

PCl₃+ 3 C₆H₅OH → P(OC₆H₅)₃+ 3 HCl

Similar reactions occur forphosphorus oxychloride,affordingtriphenylphosphate:

OPCl₃+ 3 C₆H₅OH → OP(OC₆H₅)₃+ 3 HCl

History

Etymology

The namePhosphorusin Ancient Greece was the name for the planetVenusand is derived from theGreekwords (φῶς = light, φέρω = carry), which roughly translates as light-bringer or light carrier.(InGreek mythologyand 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 withPhosphorus-the-morning-star).

According to the Oxford English Dictionary, the correct spelling of the element isphosphorus.The wordphosphorousis the adjectival form of the P³⁺valence: so, just assulfurformssulfurousandsulfuriccompounds, phosphorus forms phosphorous compounds (e.g.,phosphorous acid) and P⁵⁺valence phosphoric compounds (e.g.,phosphoric acids and phosphates).

Discovery

Robert Boyle

The discovery of phosphorus, the first element to be discovered that was not known since ancient times,is credited to the German alchemistHennig Brandin 1669, although others might have discovered phosphorus around the same time.Brand experimented withurine,which contains considerable quantities of dissolved phosphates from normal metabolism.Working inHamburg,Brand attempted to create the fabledphilosopher's stonethrough thedistillationof somesaltsby evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. It was namedphosphorus mirabilis( "miraculous bearer of light" ).

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,(NH
₄)NaHPO
₄.While the quantities were essentially correct (it took about 1,100 litres [290 US gal] of urine to make about 60 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.

Brand at first tried to keep the method secret,but later sold the recipe for 200 thalers to Johann Daniel Kraft(de)from Dresden.Krafft toured much of Europe with it, including England, where he met withRobert Boyle.The secret—that the substance was made from urine—leaked out, andJohann 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-Hanckwitz.Godfrey later made a business of the manufacture of phosphorus.

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

4NaPO
₃+ 2SiO
₂+ 10 C → 2Na
₂SiO
₃+ 10 CO +P
₄

Robert Boyle was the first to use phosphorus to ignite sulfur-tipped wooden splints, forerunners of our modern matches, in 1680.

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".

Bone ash and guano

Guanomining in the CentralChincha Islands,c. 1860

Antoine Lavoisierrecognized phosphorus as an element in 1777 afterJohan Gottlieb GahnandCarl Wilhelm Scheele,in 1769, showed thatcalcium phosphate(Ca
₃(PO
₄)
₂) is found in bones by obtaining elemental phosphorus frombone ash.

Bone ash was the major source of phosphorus until the 1840s. The method started by roasting bones, then employed the use offire clayretortsencased in a very hot brick furnace to distill out the highly toxic elemental phosphorus product.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 orcharcoalin an iron pot, and distilling off phosphorus vapour in aretort.Carbon monoxideand other flammable gases produced during the reduction process were burnt off in aflare stack.

In the 1840s, world phosphate production turned to the mining of tropical island deposits formed from bird and batguano(see alsoGuano Islands Act). These became an important source of phosphates for fertiliser in the latter half of the 19th century.

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 byJames Burgess Readmanin 1888(patented 1889),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 thematchindustry. This used bone ash for a phosphate source, as described above. The bone-ash process became obsolete when thesubmerged-arc furnace for phosphorus productionwas introduced to reduce phosphate rock.The electric furnace method allowed production to increase to the point where phosphorus could be used in weapons of war.InWorld War I,it was used in incendiaries,smoke screensand tracer bullets.A special incendiary bullet was developed to shoot athydrogen-filledZeppelinsover Britain (hydrogen being highlyflammable).DuringWorld War II,Molotov cocktailsmade of phosphorus dissolved inpetrolwere 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.

Early matches used white phosphorus in their composition, which was dangerous due to its toxicity. Murders, suicides and accidentalpoisoningsresulted 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).In addition, exposure to the vapours gave match workers a severenecrosisof 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 theBerne Convention (1906),requiring its adoption as a safer alternative for match manufacture.The toxicity of white phosphorus led to discontinuation of its use in matches.The Allies used phosphorusincendiary bombsinWorld War IIto destroy Hamburg, the place where the "miraculous bearer of light" was first discovered.

Production

Mining of phosphate rock inNauru

In 2017, the USGS estimated 68 billion tons of world reserves, where reserve figures refer to the amount assumed recoverable at current market prices; 0.261 billion tons were mined in 2016.Critical to contemporary agriculture, its annual demand is rising nearly twice as fast as the growth of the human population.The production of phosphorus may have peaked before 2011 and some scientists predict reserves will be depleted before the end of the 21st century.Phosphorus comprises about 0.1% by mass of the average rock, and consequently, the Earth's supply is vast, though dilute.

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 givephosphoric acid.Phosphoric acid is then neutralized to give various phosphate salts, which comprise fertilizers. In the wet process, phosphorus does not undergo redox.About five tons ofphosphogypsumwaste are generated per ton of phosphoric acid production. Annually, the estimated generation of phosphogypsum worldwide is 100 to 280 Mt.

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 asubmerged-arc furnacewhich is energy intensive.Presently, about 1,000,000short tons(910,000t) 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 mostlySiO
₂,andcoketo produceP
₄.TheP
₄product, being volatile, is readily isolated:

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 of Fe₂P, resulting from iron impurities in the mineral precursors. The silicateslagis 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.

Historical routes

Historically, before the development of mineral-based extractions, white phosphorus was isolated on an industrial scale frombone ash.In this process, thetricalcium phosphatein bone ash is converted tomonocalcium phosphatewithsulfuric acid:

Ca₃(PO₄)₂+ 2 H₂SO₄→ Ca(H₂PO₄)₂+ 2 CaSO₄

Monocalcium phosphate is then dehydrated to the corresponding metaphosphate:

Ca(H₂PO₄)₂→ Ca(PO₃)₂+ 2 H₂O

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

3 Ca(PO₃)₂+ 10 C → Ca₃(PO₄)₂+ 10 CO + P₄

Peak phosphorus

Annual global phosphate rock production (megatonnes per yr), 1994–2022 (data from US Geological Survey)

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 commercialraw material.The term is used in an equivalent way to the better-known termpeak oil.The issue was raised as a debate on whether phosphorus shortages might be imminent around 2010, which was largely dismissed afterUSGSand other organizationsincreased world estimates on available phosphorus resources, mostly in the form of additional resources inMorocco.However, exact reserve quantities remain uncertain, as do the possible impacts of increased phosphate use on future generations.This is important becauserock phosphateis a key ingredient in many inorganicfertilizers.Hence, a shortage in rock phosphate (or just significant price increases) might negatively affect the world'sfood security.

Phosphorus is a finite (limited) resource that is widespread in the Earth's crust and in living organisms but is relativelyscarcein concentrated forms, which are not evenly distributed across the Earth. The only cost-effective production method to date is theminingofphosphate rock,but only a few countries have significant commercialreserves.The top five areMorocco(including reserves located inWestern Sahara),China,Egypt,AlgeriaandSyria.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.

Means of commercial phosphorus production besides mining are few because thephosphorus cycledoes not include significant gas-phase transport.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.Others suggest that supplies will last for several hundreds of years.As with thetiming of peak oil,the question is not settled, and researchers in different fields regularly publish different estimates of the rock phosphate reserves.

Background

Phosphate rock mined in the United States, 1900–2015 (data from US Geological Survey)

The peak phosphorus concept is connected with the concept ofplanetary boundaries.Phosphorus, as part ofbiogeochemicalprocesses, 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.

Estimates of world phosphate reserves

Global distribution of commercial reserves of rock phosphate in 2016

The accurate determination of peakphosphorusis dependent on knowing the total world's commercialphosphatereserves and resources, especially in the form ofphosphate 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.

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,and vegetation 0.03–0.2%.Although quadrillions of tons of phosphorus exist in the Earth's crust,these are currently not economically extractable.

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.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.Phosphorus reserve figures are intensely debated.Gilbert suggest that there has been little external verification of the estimate.A 2014 reviewconcluded 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):Morocco50,China3.2,Egypt2.8,Algeria2.2,Syria1.8,Brazil1.6,Saudi Arabia1.4,South Africa1.4,Australia1.1,United States1.0,Finland1.0,Russia0.6,Jordan0.8.

Rock phosphate shortages (or just significant price increases) might negatively affect the world'sfood security.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.

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.

United States

Further information: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 ofChina.In 2019, the US produced 10 percent of the world's phosphate rock.

Exhaustion of guano reserves

Main article:Guano

In 1609Garcilaso de la Vegawrote the bookComentarios Realesin 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.In the early 1800sAlexander von Humboldtintroducedguanoas a source ofagriculturalfertilizerto Europe after having discovered it 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 metersdeep.The guano had previously been used by theMochepeople as a source of fertilizer by mining it and transporting it back toPeruby boat. International commerce in guano did not start until after 1840.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.

Phosphorus conservation and recycling

Phosphate mine onNauru,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 ofopen defecationwhich is still widespread on a global scale) or in rivers or the ocean viasewage systemsandsewage treatment plantsin the case of cities connected to sewer systems. An example of one crop that takes up large amounts of phosphorus issoy.

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 insanitationsystems. TheSoil 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.One potential solution to the shortage of phosphorus is greater recycling of human and animal wastes back into the environment.

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 asno-till farming,terracing,contour tilling,and the use ofwindbreakshave 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.

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 article:Excreta reuse

The oldest method of recycling phosphorus is through the reuse of animalmanureand humanexcretain 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 theGlobal Souththese processes are not always followed, leading to outbreaks of diseases transmitted via thefecal–oral routesuch ascholera.

Sewage sludge

Sewage treatment plants that have anenhanced biological phosphorus removalstep produce asewage sludgethat is rich in phosphorus. Various processes have been developed to extract phosphorus from sewage sludge directly, from the ash afterincinerationof the sewage sludge or from other products ofsewage sludge treatment.This includes the extraction of phosphorus rich materials such asstruvitefrom waste processing plants.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.

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.

Neutron transmutation doping

The above routes refer to "production" in thechemicalsense 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 anuclearsense 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.

Neutrontransmutationdoping (NTD) is an unusual doping method for special applications. Most commonly, it is used to dope silicon n-type in high-power electronics andsemiconductor detectors.It is based on the conversion of the³⁰Si isotope into phosphorus atoms byneutron absorptionandbeta decayas follows:

$$^{30}\mathrm {Si} \,(n,\gamma )\,^{31}\mathrm {Si} \rightarrow \,^{31}\mathrm {P} +\beta ^{-}\;(T_{1/2}=2.62\mathrm {h} ).$$ In practice, the silicon is typically placed near or inside anuclear reactor(most commonly aresearch reactore.g. the one atMIT) 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.

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.

Food additive

Phosphorus is an essentialmineralfor humans listed in theDietary Reference Intake(DRI).

Food-gradephosphoric acid(additiveE338) is used to acidify foods and beverages such as variouscolasand jams, providing a tangy or sour taste. The phosphoric acid also serves as apreservative.Soft drinks containing phosphoric acid, includingCoca-Cola,are sometimes calledphosphate sodasor phosphates. Phosphoric acid in soft drinks has the potential to cause dental erosion.Phosphoric acid also has the potential to contribute to the formation ofkidney stones,especially in those who have had kidney stones previously.

Fertiliser

Main article:Fertiliser

See also:Phosphorus cycle

Phosphorus is an essential plant nutrient (the most often limiting nutrient, afternitrogen),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.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.Heavy use of phosphorus fertilizers and their runoff have resulted ineutrophication(overenrichment) ofaquatic ecosystems.

Natural phosphorus-bearing compounds are mostly inaccessible to plants because of the low solubility and mobility in soil.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 cycleis 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.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 acidand the greatest industrial use of elementalsulfur.

Widely used compounds	Use
Ca(H2PO4)2·H2O	Baking powder and fertilisers
CaHPO4·2H2O	Animal food additive, toothpowder
H3PO4	Manufacture of phosphate fertilisers
PCl3	Manufacture of POCl3and pesticides
POCl3	Manufacture of plasticiser
P4S10	Manufacturing of additives and pesticides
Na5P3O10	Detergents

Organophosphorus

White phosphorus is widely used to makeorganophosphorus compoundsthrough intermediatephosphorus chloridesand two phosphorus sulfides,phosphorus pentasulfideandphosphorus sesquisulfide.Organophosphorus compounds have many applications, including inplasticisers,flame retardants,pesticides,extraction agents, nerve agents andwater treatment.

Metallurgical aspects

Phosphorus is also an important component insteelproduction, in the making ofphosphor bronze,and in many other related products.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 embrittlementresistance than normal copper. Phosphate conversion coatingis a chemical treatment applied to steel parts to improve their corrosion resistance.

Matches

Match striking surface made of a mixture of red phosphorus, glue and ground glass. The glass powder is used to increase the friction.

Main article:Match

The first striking match with a phosphorus head was invented byCharles Sauriain 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,sensitive to storage conditions, toxic if ingested, and hazardous when accidentally ignited on a rough surface.Production in several countries was banned between 1872 and 1925.The internationalBerne 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 byphosphorus 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 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 headpotassium chlorate,an oxygen-releasing compound. When struck, small amounts ofabrasionfrom 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. Nonetheless, safety matches, invented in 1844 byGustaf Erik Paschand market ready by the 1860s, did not gain consumer acceptance until the prohibition of white phosphorus. Using a dedicated striker strip was considered clumsy.

Water softening

Sodium tripolyphosphatemade from phosphoric acid is used in laundry detergents in some countries, but banned for this use in others.This compound softens the water to enhance the performance of the detergents and to prevent pipe/boiler tubecorrosion.

Miscellaneous

• Phosphates are used to make special glasses forsodium lamps.
• Bone-ash (mostlycalcium phosphate) is used in the production of fine china.
• Phosphoric acid made from elemental phosphorus is used in food applications such assoft drinks,and as a starting point for food grade phosphates.These includemonocalcium phosphateforbaking powderandsodium tripolyphosphate.Phosphates are used to improve the characteristics of processed meat andcheese,and in toothpaste.
• White phosphorus,called "WP" (slang term "Willie Peter" ) is used inmilitaryapplications asincendiary bombs,forsmoke-screeningas smoke pots andsmoke bombs,and intracer ammunition.It is also a part of an obsoleteM34 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.Military uses of white phosphorus are constrained by international law.
• ³²P and³³P are used as radioactive tracers in biochemical laboratories.
• Phosphorus is adopantinn-type semiconductors

Biological role

See also:Calcium metabolism

Inorganic phosphorus in the form of the phosphatePO³⁻
₄is required for all known forms oflife.Phosphorus plays a major role in the structural framework ofDNAandRNA.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.Phospholipidsare the main structural components of all cellular membranes.Calcium phosphatesalts assist in stiffeningbones.Biochemists commonly use the abbreviation "P_i"to refer to inorganic phosphate.

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 fromglycerolwith 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.

An average adult human contains about 0.7 kg 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. Average phosphorus concentration in the blood is about 0.4 g/L; about 70% of that is organic and 30% inorganic phosphates.An adult with healthy diet consumes and excretes about 1–3 grams of phosphorus per day, with consumption in the form of inorganic phosphate and phosphorus-containing biomolecules such asnucleic acidsandphospholipids;and excretion almost exclusively in the form of phosphate ions such asH
₂PO⁻
₄andHPO²⁻
₄.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 ishydroxyapatiteas well as amorphous forms of calcium phosphate, possibly including carbonate. Hydroxyapatite is the main component of tooth enamel.Water fluoridationenhances the resistance of teeth to decay by the partial conversion of this mineral to the still harder materialfluorapatite:

Ca
₅(PO
₄)
₃OH+F⁻
→Ca
₅(PO
₄)
₃F+OH⁻

Phosphorus deficiency

In medicine, phosphate deficiency syndrome may be caused bymalnutrition,by failure to absorb phosphate, and by metabolic syndromes that draw phosphate from the blood (such as inrefeeding syndromeafter malnutrition) 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.

Phosphorus is an essentialmacromineralfor plants, which is studied extensively inedaphologyto understand plant uptake fromsoilsystems. Phosphorus is alimiting factorin manyecosystems;that is, the scarcity of phosphorus limits the rate of organism growth. An excess of phosphorus can also be problematic, especially in aquatic systems whereeutrophicationsometimes leads toalgal blooms.

Nutrition

Dietary recommendations

TheU.S. Institute of Medicine(IOM) updated Estimated Average Requirements (EARs) andRecommended Dietary Allowances(RDAs) for phosphorus in 1997. If there is not sufficient information to establish EARs and RDAs, an estimate designatedAdequate Intake(AI) is used instead. The current EAR 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 levels(ULs) for vitamins and minerals when evidence is sufficient. In the case of phosphorus, the UL is 4000 mg/day. Collectively the EARs, RDAs, AIs and ULs are referred to asDietary Reference Intakes(DRIs).

TheEuropean Food Safety Authority(EFSA) refers to the collective set of information as Dietary Reference Values, withPopulation Reference Intake(PRI) instead of RDA, and Average Requirement instead of EAR. 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.The European Food Safety Authority reviewed the same safety question and decided that there was not sufficient information to set a UL.

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 mg, but as of May 27, 2016, it was revised to 1250 mg to bring it into agreement with the RDA.A table of the old and new adult daily values is provided atReference Daily Intake.

Food sources

The main food sources for phosphorus are the same as those containingprotein,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.

Precautions

Phosphorus explosion

Organic compounds of phosphorus form a wide class of materials; many are required for life, but some are extremely toxic. Fluorophosphateestersare among the most potentneurotoxinsknown. A wide range of organophosphorus compounds are used for their toxicity aspesticides(herbicides,insecticides,fungicides,etc.) andweaponisedas nerve agents against enemy humans. Most inorganic phosphates are relatively nontoxic and essential nutrients.

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 istoxic,causing severe liver damage on ingestion and may cause a condition known as "Smoking Stool Syndrome".

In the past, external exposure to elemental phosphorus was treated by washing the affected area with 2%copper(II) sulfatesolution to form harmless compounds that are then washed away. According to the recentUS 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."

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. Promptlydebridethe 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-basedointmentsuntil it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns. "As white phosphorus readily mixes with oils, any oily substances or ointments are not recommended until the area is thoroughly cleaned and all white phosphorus removed.

People can be exposed to phosphorus in the workplace by inhalation, ingestion, skin contact, and eye contact. TheOccupational Safety and Health Administration(OSHA) has set the phosphorus exposure limit (Permissible exposure limit) in the workplace at 0.1 mg/m³over an 8-hour workday. TheNational Institute for Occupational Safety and Health(NIOSH) has set aRecommended exposure limit(REL) of 0.1 mg/m³over an 8-hour workday. At levels of 5 mg/m³,phosphorus isimmediately dangerous to life and health.

US DEA List I status

Phosphorus can reduce elementaliodinetohydroiodic acid,which is a reagent effective for reducingephedrineorpseudoephedrinetomethamphetamine.For this reason, red and white phosphorus were designated by the United StatesDrug Enforcement AdministrationasList I precursor chemicalsunder21 CFR 1310.02effective on November 17, 2001.In the United States, handlers of red or white phosphorus are subject to stringent regulatory controls.

See also

• Hubbert peak theory

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• Ruttenberg, K. C.Phosphorus Cycle – Terrestrial Phosphorus Cycle, Transport of Phosphorus,fromContinents to the Ocean, The Marine Phosphorus Cycle.ArchivedJuly 13, 2011, atarchive.today.
• Sommers, Michael A. (2007-08-15).Phosphorus.Rosen Group.ISBN978-1-4042-1960-1.{{cite book}}:CS1 maint: date and year (link)
• Schrödter, Klaus; Bettermann, Gerhard; Staffel, Thomas; Wahl, Friedrich; Klein, Thomas; Hofmann, Thomas. "Phosphoric Acid and Phosphates".Ullmann's Encyclopedia of Industrial Chemistry.Weinheim: Wiley-VCH.doi:10.1002/14356007.a19_465.pub3.ISBN978-3527306732.
• Threlfall, Richard E. (1951).The Story of 100 years of Phosphorus Making: 1851–1951.Oldbury: Albright & Wilson Ltd.
• Toy, Arthur D. F. (1975).The Chemistry of Phosphorus.Texts in Inorganic Chemistry. Vol. 3. Pergamon.ISBN978-1-4831-4741-3.Retrieved2013-10-22.

Further reading

• Podger, Hugh (2002).Albright & Wilson. The Last 50 years.Studley: Brewin Books.ISBN1-85858-223-7.
• 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.)