Phosphorus

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

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 (see Allotropes of phosphorus)
Appearance	white, red and violet are waxy, black is metallic-looking
Standard atomic weight Ar°(P)
	30.973761998±0.000000005 [1] 30.974±0.001 (abridged) [2]
Abundance
in the  Earth's crust	5.2 (silicon = 100)
Phosphorus in the periodic 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 ] 3s2 3p3
Electrons per shell	2, 8, 5
Physical properties
Phase at  STP	solid
Melting point	white: 317.3  K ​(44.15 °C,​111.5 °F) red: ∼860 K (∼590 °C, ∼1090 °F) [3]
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 (near  r.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: 0.66  kJ/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 at T (K) 279 307 342 388 453 549
Vapour pressure(red) P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 455 489 529 576 635 704
Atomic properties
Oxidation states	common: −3, +3, +5 −2, [4] −1, [4] 0, [5] +1, [4] [6] +2, [4] +4 [4]
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±3  pm
Van der Waals radius	180 pm
Spectral lines of phosphorus
Other properties
Natural occurrence	primordial
Crystal structure	α-white:​ body-centred cubic (bcc)(cI232)
Lattice constant	a = 1.869 nm (at 20 °C) [7]
Crystal structure	black:​ orthorhombic (oS8)
Lattice constants	a = 0.33137 nm b = 1.0477 nm c = 0.43755 nm (at 20 °C) [7]
Thermal conductivity	white: 0.236 W/(m⋅K) black: 12.1 W/(m⋅K)
Magnetic ordering	white, red, violet, black: diamagnetic [8]
Molar magnetic susceptibility	−20.8×10−6 cm3/mol(293 K) [9]
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 [10] (1777)
Isotopes of phosphorus v e
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
Category: Phosphorus view talk edit | references

Phosphorus is a chemical element with the symbol P and the atomic number 15. All elemental forms of phosphorus are highly reactive and are therefore never found in nature. They can nevertheless be prepared artificially, the two most common ones being white phosphorus and red phosphorus. With ³¹P as its only stable isotope, phosphorus has an occurrence in Earth's crust of about 0.1%, generally as phosphate rock. A member of the pnictogen family, phosphorus readily forms a wide variety of organic and inorganic compounds, with as its main oxidation states +5, +3 and -3.

The isolation of white phosphorus in 1669 by Hennig Brand marked the first "discovery" of an element not known since Antiquity. The name phosphorus is a reference to the god of the Morning star in Greek mythology, inspired by the faint glow of white phosphorus when exposed to oxygen. This property is also at the origin of the term phosphorescence , meaning glow after illumination, although white phosphorus itself does not exhibit phosphorescence, but chemiluminescence caused by its oxidation. Its high toxicity makes exposure to white phosphorus very dangerous, while its flammability and pyrophoricity can be weaponized in the form of incendiaries. Red phosphorus is less dangerous and is used in matches and fire retardants.

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

Phosphorus is essential to all known forms of life, largely through organophosphates, organic compounds containing the phosphate ion PO3−4 as a functional group. These include DNA, RNA, ATP, and phospholipids, complex compounds fundamental to the functioning of all cells. The main component of bones and teeth, bone mineral, is a modified form of hydroxyapatite, itself a phosphorus mineral.

History

Discovery

The Alchemist in Search of the Philosophers Stone (1771), by Joseph Wright, depicting Hennig Brand discovering phosphorus.

Phosphorus was the first element to be "discovered", in the sense that it was not known since ancient times. ^[11] The discovery is credited to the Hamburg alchemist Hennig Brand in 1669, who was attempting to create the fabled philosopher's stone. ^[12] To this end, he experimented with urine, 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, then distilling it at a high temperature and leading the resulting vapours through water, he obtained a white, waxy substance that glowed in the dark and burned brilliantly. He named it in Latin : 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 from Greek mythology, where it references the god of the morning star, also known as the planet Venus. ^{[13] [15]}

Brand at first tried to keep the method secret, ^[16] but later sold the recipe for 200 thalers to Johann Daniel Kraft  [ de ] from Dresden. ^[13] Kraft toured much of Europe with it, including London, where he met with Robert Boyle. The crucial fact that the substance was made from urine was eventually found out, and Johann 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 ignite sulfur-tipped wooden splints, forerunners of modern matches, ^[17] and also improved the process by using sand in the reaction:

4 NaPO₃ + 2 SiO₂ + 10 C → 2 Na₂SiO₃ + 10 CO + P₄

Boyle's assistant Ambrose Godfrey-Hanckwitz later made a business of the manufacture of phosphorus.

Bone ash

Antoine Lavoisier recognized phosphorus as an element in 1777 after Johan Gottlieb Gahn and Carl Wilhelm Scheele showed in 1769 that calcium phosphate is found in bones by obtaining elemental phosphorus from bone ash. ^[10] Bone ash subsequently became the primary industrial source of phosphorus and remained so until the 1840s. ^[18] The process consisted of several steps. ^{[19] [20]} First, grinding up the bones into their constituent tricalcium phosphate and treating it with sulfuric acid:

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

Then, dehydrating the resulting monocalcium phosphate:

Ca(H₂PO₄)₂ → Ca(PO₃)₂ + 2 H₂O

Finally, mixing the obtained calcium metaphosphate with ground coal or charcoal in an iron pot, and distilling phosphorus vapour out of a retort:

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

This way, two-thirds of the phosphorus was turned into white phosphorus while one-third remained in the residue as calcium orthophosphate. The carbon monoxide produced during the reaction process was burnt off in a flare stack.

Guano

Main article: Guano

Guano mining in the Central Chincha Islands, c. 1860

In 1609 Inca Garcilaso de la Vega wrote the book Comentarios Reales in which he described many of the agricultural practices of the Incas prior to the arrival of the Spaniards and introduced the use of guano as a fertiliser. As Garcilaso described, the Incas near the coast harvested guano. ^[21] In the early 1800s Alexander von Humboldt introduced guano as a source of agricultural fertiliser to Europe after having discovered it in exploitable quantities on islands off the coast of South America. It has been reported that, at the time of its discovery, the guano on some islands was over 30 meters deep. ^[22] The guano had previously been used by the Moche people as a source of fertiliser by mining it and transporting it back to Peru by boat. International commerce in guano did not start until after 1840. ^[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 of superphosphate.

Matches

Further information: Match § History

Early matches used white phosphorus in their composition, which was dangerous due to its toxicity. Exposure to the vapours gave match workers a severe necrosis of the bones of the jaw, known as "phossy jaw". When a safe process for manufacturing red phosphorus was discovered, with its far lower flammability and toxicity, laws were enacted, under the Berne Convention (1906), requiring its adoption as a safer alternative for match manufacture. ^[23] The toxicity of white phosphorus led to discontinuation of its use in matches. ^[24]

Phosphate rock

A worker tends an electric phosphate smelting furnace in Muscle Shoals, Alabama, 1942

Phosphate rock, which usually contains calcium phosphate, was first used in 1850 to make phosphorus. With the introduction of the submerged-arc furnace for phosphorus production by James Burgess Readman in 1888 ^[25] (patented 1889), ^[26] the use of bone-ash became obsolete. ^{[27] [28]} 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

Further information: White phosphorus munition § History

White phosphorus shell explosion in France during the First World War (1918)

The electric furnace method allowed production to increase to the point where white phosphorus could be used in weapons of war. In World War I, it was used in incendiary ammunition, smoke screens and tracer ammunition. A special incendiary bullet was developed to shoot at hydrogen-filled Zeppelins over Britain (hydrogen being highly flammable). ^[29]

During World War II, Molotov cocktails made of phosphorus dissolved in petrol were distributed in Britain to specially selected civilians as part of the preparations for a potential invasion. The United States also developed the M15 white-phosphorus hand grenade, a precursor to the M34 grenade, while the British introduced the similar No 77 grenade. These multipurpose grenades were mostly used for signaling and smoke screens, although they were also efficient anti-personnel weapons. ^[30] The difficulty to extinguish burning phosphorus and the very severe burns it causes had a strong psychological impact on the enemy. ^[31] Phosphorus incendiary bombs were used on a large scale, notably to destroy Hamburg, the place where the "miraculous bearer of light" was first discovered. ^[15]

Characteristics

Isotopes

Main article: Isotopes of phosphorus

There are 22 known isotopes of phosphorus, ^[32] ranging from ²⁶P to ⁴⁷P. ^[33] Only ³¹P is stable and is therefore present at 100% abundance. The half-integer nuclear spin and high abundance of ³¹P make phosphorus-31 NMR spectroscopy a very useful analytical tool in studies of phosphorus-containing samples.

Two radioactive isotopes of phosphorus have half-lives suitable for biological scientific experiments, and are used as radioactive tracers in biochemical laboratories. ^[34] These are:

• ³²P , a beta-emitter (1.71 MeV) with a half-life of 14.3 days, which is used routinely in life-science laboratories, primarily to produce radiolabeled DNA and RNA probes, e.g. for use in Northern blots or Southern blots.
• 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 as DNA sequencing.

The high-energy beta particles from ³²P penetrate skin and corneas and any ³²P ingested, inhaled, or absorbed is readily incorporated into bone and nucleic acids. For these reasons, personnel working with ³²P 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 of X-rays via Bremsstrahlung (braking radiation) in dense shielding materials such as lead. Therefore, the radiation must be shielded with low density materials such as water, acrylic or other plastic. ^[35]

Allotropes

Main article: Allotropes of phosphorus

Crystalline structures of the main phosphorus allotropes

White

Red

Violet

Black

Phosphorus has several allotropes that exhibit very diverse properties. ^[36] The most useful and therefore common is white phosphorus, followed by red phosphorus. The two other main allotropes, violet and black phosphorus, have either a more fundamental interest or specialized applications. ^[37] Many other allotropes have been theorized and synthesized, with the search for new materials an active area of research. ^[38] 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

White phosphorus exposed to air glows in the dark.

White phosphorus is a soft, waxy molecular solid that is insoluble in water. ^[31] It is also very toxic, highly flammable and pyrophoric, igniting in air at about 30 °C (303 K). ^[39] Structurally, it is composed of P₄ tetrahedra. The nature of bonding in a given P₄ tetrahedron can be described by spherical aromaticity or cluster bonding, that is the electrons are highly delocalized. This has been illustrated by calculations of the magnetically induced currents, which sum up to 29 nA/T, much more than in the archetypical aromatic molecule benzene (11 nA/T). ^[40] The P₄ molecule in the gas phase has a P-P bond length of 2.1994(3) Å as determined by gas electron diffraction. ^[40]

White phosphorus exists in two crystalline forms named α (alpha) and β (beta), differing in terms of the relative orientation of the constituent P₄ tetrahedra. ^{[41] [42]} The α-form is most stable at room temperature and has a cubic crystal structure. When cooled down to 195.2 K (−78.0 °C) it transforms into the β-form, turning into an hexagonal 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 facing thermal decomposition at 1,100 K (830 °C) where it turns into gaseous diphosphorus (P₂). ^[43] This molecule 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. ^[44] At still higher temperatures, P₂ dissociates into atomic P. ^[31]

When exposed to air, white phosphorus faintly glows green and blue due to oxidation, a phenomenon best visible in the dark. This reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived molecules HPO and P₂O₂ that both emit visible light. ^[45] However, in a pure-oxygen environment phosphorus does not glow at all, with the oxidation happening only in a range of partial pressures. ^[46] Derived from this phenomenon, the terms phosphors and phosphorescence have been loosely used to describe substances that shine in the dark. However, phosphorus itself is not phosphorescent but chemiluminescent, since it glows due to a chemical reaction and not the progressive reemission of previously absorbed light. ^[14]

Red phosphorus

Red phosphorus is polymeric in structure. It can be viewed as a derivative of P₄ wherein one P-P bond is broken and one additional bond is formed with the neighbouring tetrahedron, resulting in chains of P₂₁ molecules linked by van der Waals forces. ^[47] 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 is amorphous, but can be crystallised upon further heating into violet phosphorus or fibrous red phosphorus depending on the reaction conditions. Red phosphorus is therefore not an allotrope in the strictest sense of the term, but rather an intermediate between other crystalline allotropes of phosphorus, and consequently most of its properties have a range of values. Freshly prepared, bright red phosphorus is highly reactive and ignites at about 300 °C (573 K). ^[48] After prolonged heating or storage, the color darkens; the resulting product is more stable and does not spontaneously ignite in air. ^[49]

Violet phosphorus

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 molten lead, a red/purple form is obtained. Therefore, this form is sometimes known as "Hittorf's phosphorus" . ^[50]

Black phosphorus

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 resembles graphite, being black and flaky, a conductor of electricity, and having puckered sheets of linked atoms. ^{[51] [52] [53]} 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. ^[54] Single-layer black phosphorus is called phosphorene, and is therefore predictably analogous to graphene.

Compounds

Main category: Phosphorus compounds

Phosphorus(V)

The tetrahedral structure of P₄O₁₀ and P₄S₁₀

The most prevalent compounds of phosphorus are derivatives of phosphate (PO3−4), a tetrahedral anion. ^[55] 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 (K_a1 = 7.25×10⁻³)

H₂PO−4 + H₂O ⇌ H₃O⁺ + HPO2−4 (K_a2 = 6.31×10⁻⁸)

HPO2−4 + H₂O ⇌ H₃O⁺ + PO3−4 (K_a3 = 3.98×10⁻¹³)

Phosphate exhibits a tendency to form chains and rings containing P-O-P bonds. Many polyphosphates are known, including ATP. Polyphosphates arise by dehydration of hydrogen phosphates such as HPO2−4 and H₂PO−4. For example, the industrially important pentasodium triphosphate (also known as sodium tripolyphosphate, STPP) is produced industrially by the megatonne by this condensation reaction:

2 Na₂HPO₄ + NaH₂PO₄ → Na₅P₃O₁₀ + 2 H₂O

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

With metal cations, 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).

Phosphorus halides such as PF₅ and PCl₅ are common compounds. PF₅ is a colourless gas and the molecules have trigonal bipyramidal geometry. PCl₅ is a colourless solid which has an ionic formulation of PCl+4PCl−6, but adopts the trigonal bipyramidal geometry when molten or in the vapour phase. ^[31] PBr₅ is an unstable solid formulated as PBr+4Br⁻ and PI₅ is not known. ^[31] The pentachloride and pentafluoride are Lewis acids. With fluoride, PF₅ forms PF−6, an anion that is isoelectronic with SF₆. The most important oxyhalide is phosphorus oxychloride, (POCl₃), which is approximately tetrahedral.

Before extensive computer calculations were feasible, it was thought that bonding in phosphorus(V) compounds involved d orbitals. Computer modeling of molecular orbital theory indicates that this bonding involves only s- and p-orbitals. ^[56]

Phosphorus(III)

All four symmetrical trihalides are well known: gaseous PF₃ , the yellowish liquids PCl₃ and PBr₃ , and the solid PI₃ . These materials are moisture sensitive, hydrolysing to give phosphorous 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. PF₃ is toxic because it binds to haemoglobin.

Phosphorus(III) oxide, P₄O₆ (also called tetraphosphorus hexoxide) is the anhydride of P(OH)₃, the minor tautomer of phosphorous acid. The structure of P₄O₆ is like that of P₄O₁₀ without the terminal oxide groups.

Symmetric phosphorus(III) trithioesters (e.g. P(SMe)₃) can be produced from the reaction of white phosphorus and the corresponding disulfide, or phosphorus(III) halides and thiolates. Unlike the corresponding esters, they do not undergo a variant of the Michaelis-Arbuzov reaction with electrophiles, instead reverting to another phosphorus(III) compound through a sulfonium intermediate. ^[57]

Phosphorus(I) and phosphorus(II)

A stable diphosphene, a derivative of phosphorus(I)

These compounds generally feature P–P bonds. ^[31] 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

Phosphides arise by reaction of metals with red phosphorus. The alkali metals (group 1) and alkaline earth metals can form ionic compounds containing the phosphide ion, P−3. These compounds react with water to form phosphine. Other phosphides, 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. ^[31] Schreibersite is a naturally occurring metal-rich phosphide found in meteorites. The structures of the metal-rich and phosphorus-rich phosphides can be complex.

Phosphine (PH₃) and its organic derivatives 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 of calcium 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. ^[31] The highly flammable gas diphosphine (P₂H₄) is an analogue of hydrazine.

Oxoacids

Phosphorus oxoacids 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. ^[31] Although many oxoacids of phosphorus are formed, only nine are commercially important, and three of them, hypophosphorous acid, phosphorous acid, and phosphoric acid, are particularly important.

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 crystalline phosphorus nitride decomposition at 1100 K. Similarly, H₂PN is considered unstable, and phosphorus nitride halogens like F₂PN, Cl₂PN, Br₂PN, and I₂PN oligomerise into cyclic polyphosphazenes. For example, compounds of the formula (PNCl₂)_n exist mainly as rings such as the trimer hexachlorophosphazene. 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 by alkoxide (RO⁻), a family of polymers is produced with potentially useful properties. ^[58]

Sulfides

Main article: phosphorus sulfides

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 P₄S₃ is used in strike-anywhere matches. P₄S₁₀ and P₄O₁₀ have analogous structures. ^[59] 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 PCl₃ serves as a source of P³⁺ in routes to organophosphorus(III) compounds. For example, it is the precursor to triphenylphosphine:

PCl₃ + 6 Na + 3 C₆H₅Cl → P(C₆H₅)₃ + 6 NaCl

Treatment of phosphorus trihalides with alcohols and phenols gives phosphites, e.g. triphenylphosphite:

PCl₃ + 3 C₆H₅OH → P(OC₆H₅)₃ + 3 HCl

Similar reactions occur for phosphorus oxychloride, affording triphenylphosphate:

OPCl₃ + 3 C₆H₅OH → OP(OC₆H₅)₃ + 3 HCl

Occurrence

Universe

In 2013, astronomers detected phosphorus in Cassiopeia A, which confirmed that this element is produced in supernovae as a byproduct of supernova nucleosynthesis. The phosphorus-to-iron ratio in material from the supernova remnant could be up to 100 times higher than in the Milky Way in general. ^[60]

In 2020, astronomers analysed ALMA and ROSINA data from the massive star-forming region AFGL 5142, to detect phosphorus-bearing molecules and how they are carried in comets to the early Earth. ^[61]

Earth's Crust

Phosphorus has a concentration in the Earth's crust of about one gram per kilogram (compare copper at about 0.06 grams). It is not found free in nature, but is widely distributed in many minerals, usually as phosphates. ^[37] Inorganic phosphate rock, which is partially made of apatite, is today the chief commercial source of this element. According to the US Geological Survey (USGS), about 50 percent of the global phosphorus reserves are in Morocco, Algeria and Tunisia. ^[62] 85% of Earth's known reserves are in Morocco with smaller deposits in China, Russia, ^[63] Florida, Idaho, Tennessee, Utah, and elsewhere. ^[64] Albright and Wilson in the UK and their Niagara Falls plant, for instance, were using phosphate rock in the 1890s and 1900s from Tennessee, Florida, and the Îles du Connétable (guano island sources of phosphate); by 1950, they were using phosphate rock mainly from Tennessee and North Africa. ^[65]

Production

Mining of phosphate rock in Nauru

Wet process

Most phosphorus-bearing material is for agriculture fertilisers. In this case where the standards of purity are modest, phosphorus is obtained from phosphate rock by what is called the "wet process." The minerals are treated with sulfuric acid to give phosphoric acid. Phosphoric acid is then neutralized to give various phosphate salts, which comprise fertilizers. In the wet process, phosphorus does not undergo redox. ^[66] About five tons of phosphogypsum waste are generated per ton of phosphoric acid production. Annually, the estimated generation of phosphogypsum worldwide is 100 to 280 Mt. ^[67]

Thermal process

For the use of phosphorus in drugs, detergents, and foodstuff, the standards of purity are high, which led to the development of the thermal process. In this process, phosphate minerals are converted to white phosphorus, which can be purified by distillation. The white phosphorus is then oxidised to phosphoric acid and subsequently neutralised with a base to give phosphate salts. The thermal process is conducted in a submerged-arc furnace which is energy intensive. ^[66] Presently, about 1,000,000 short tons (910,000  t ) of elemental phosphorus is produced annually. Calcium phosphate (as phosphate rock), mostly mined in Florida and North Africa, can be heated to 1,200–1,500 °C with sand, which is mostly SiO₂, and coke to produce P₄. The P₄ product, being volatile, is readily isolated: ^[68]

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 include ferrophosphorus, a crude form of Fe₂P, resulting from iron impurities in the mineral precursors. The silicate slag is a useful construction material. The fluoride is sometimes recovered for use in water fluoridation. More problematic is a "mud" containing significant amounts of white phosphorus. Production of white phosphorus is conducted in large facilities in part because it is energy intensive. The white phosphorus is transported in molten form. Some major accidents have occurred during transportation. ^[69]

Peak phosphorus

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

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. ^[71] Critical to contemporary agriculture, its annual demand is rising nearly twice as fast as the growth of the human population. ^[63] The production of phosphorus may have peaked before 2011 and some scientists predict reserves will be depleted before the end of the 21st century. ^{[72] [63]} Phosphorus comprises about 0.1% by mass of the average rock, and consequently, the Earth's supply is vast, though dilute. ^[31]

Peak phosphorus is a concept to describe the point in time when humanity reaches the maximum global production rate of phosphorus as an industrial and commercial raw material. The term is used in an equivalent way to the better-known term peak oil. ^[73] The issue was raised as a debate on whether phosphorus shortages might be imminent around 2010, which was largely dismissed after USGS and other organizations ^[74] increased world estimates on available phosphorus resources, mostly in the form of additional resources in Morocco. However, exact reserve quantities remain uncertain, as do the possible impacts of increased phosphate use on future generations. ^[75] This is important because rock phosphate is a key ingredient in many inorganic fertilizers. Hence, a shortage in rock phosphate (or just significant price increases) might negatively affect the world's food security. ^[76]

Phosphorus is a finite (limited) resource that is widespread in the Earth's crust and in living organisms but is relatively scarce in concentrated forms, which are not evenly distributed across the Earth. The only cost-effective production method to date is the mining of phosphate rock, but only a few countries have significant commercial reserves. The top five are Morocco (including reserves located in Western Sahara), China, Egypt, Algeria and Syria. ^[77] 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. ^[78]

Means of commercial phosphorus production besides mining are few because the phosphorus cycle does not include significant gas-phase transport. ^[79] 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. ^{[73] [72]} Others suggest that supplies will last for several hundreds of years. ^[80] As with the timing of peak oil, the question is not settled, and researchers in different fields regularly publish different estimates of the rock phosphate reserves. ^[81]

Background

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

The peak phosphorus concept is connected with the concept of planetary boundaries. Phosphorus, as part of biogeochemical processes, belongs to one of the nine "Earth system processes" which are known to have boundaries. As long as the boundaries are not crossed, they mark the "safe zone" for the planet. ^[82]

Estimates of world phosphate reserves

Global distribution of commercial reserves of rock phosphate in 2016

The accurate determination of peak phosphorus is dependent on knowing the total world's commercial phosphate reserves and resources, especially in the form of phosphate rock (a summarizing term for over 300 ores of different origin, composition, and phosphate content). "Reserves" refers to the amount assumed recoverable at current market prices and "resources" refers to estimated amounts of such a grade or quality that they have reasonable prospects for economic extraction. ^{[84] [85]}

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, ^[86] and vegetation 0.03–0.2%. ^[87] Although quadrillions of tons of phosphorus exist in the Earth's crust, ^[88] these are currently not economically extractable.

US production of phosphate rock peaked in 1980 at 54.4 million metric tons. The United States was the world's largest producer of phosphate rock from at least 1900, up until 2006, when US production was exceeded by that of China. In 2019, the US produced 10 percent of the world's phosphate rock. ^[89]

In 2023, the United States Geological Survey (USGS) estimated that economically extractable phosphate rock reserves worldwide are 72 billion tons, while world mining production in 2022 was 220 million tons. ^[77] Assuming zero growth, the reserves would thus last for around 300 years. This broadly confirms a 2010 International Fertilizer Development Center (IFDC) report that global reserves would last for several hundred years. ^{[80] [74]} Phosphorus reserve figures are intensely debated. ^{[84] [90] [91]} Gilbert suggest that there has been little external verification of the estimate. ^[92] A 2014 review ^[81] concluded that the IFDC report "presents an inflated picture of global reserves, in particular those of Morocco, where largely hypothetical and inferred resources have simply been relabeled “reserves".

The countries with most phosphate rock commercial reserves (in billion metric tons): Morocco 50, China 3.2, Egypt 2.8, Algeria 2.2, Syria 1.8, Brazil 1.6, Saudi Arabia 1.4, South Africa 1.4, Australia 1.1, United States 1.0, Finland 1.0, Russia 0.6, Jordan 0.8. ^{[93] [77] [94]}

Rock phosphate shortages (or just significant price increases) might negatively affect the world's food security. ^[76] 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. ^[95]

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. ^[96]

Conservation and recycling

A phosphate removal sewage treatment station in Yorkshire, England

Reducing agricultural runoff and soil erosion can slow the frequency with which farmers have to reapply phosphorus to their fields. Agricultural methods such as no-till farming, terracing, contour tilling, and the use of windbreaks 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. ^[97]

Sewage treatment plants that have a dedicated phosphorus removal step produce phosphate-rich sewage sludge that can then be treated to extract phosphorus from it. This is done by incinerating the sludge and recovering the resulting ash.. ^[98] Another approach lies into the recovery of phosphorus-rich materials such as struvite from waste processing plants, which is done by adding magnesium to the waste. ^[92] However, the technologies currently in use are not yet cost-effective, given the current price of phosphorus on the world market. ^{[99] [100]}

Applications

Flame retardants and plasticisers

Phosphorus compounds are used as flame retardants. ^[101] Some of these, such as Tricresyl phosphate and 2-Ethylhexyl diphenyl phosphate, are also plasticisers, making these two properties useful in the production of non-flammable plastic products and derivatives. ^{[31] [102]}

Food additive

Food-grade phosphoric acid (additive E338 ^[103] ) is used to acidify foods and beverages such as various colas and jams, providing a tangy or sour taste. ^[104] The phosphoric acid also serves as a preservative. ^[105] Soft drinks containing phosphoric acid, including Coca-Cola, are sometimes called phosphate sodas or phosphates. Phosphoric acid in soft drinks has the potential to cause dental erosion, ^[106] as well as contribute to the formation of kidney stones, especially in those who have had kidney stones previously. ^[107] Phosphates are used to improve the characteristics of processed meat and cheese, in baking powder, and in toothpaste. ^[104]

Fertiliser

Main article: Fertiliser

See also: Phosphorus cycle

Phosphorus is an essential plant nutrient (the most often limiting nutrient, after nitrogen), ^[108] and the bulk of all phosphorus production is in concentrated phosphoric acids for agriculture fertilisers, containing as much as 70% to 75% P₂O₅. That led to large increase in phosphate (PO₄³⁻) production in the second half of the 20th century. ^[63] Artificial phosphate fertilisation is necessary because phosphorus is essential to all living organisms; it is involved in energy transfers, strength of root and stems, photosynthesis, the expansion of plant roots, formation of seeds and flowers, and other important factors effecting overall plant health and genetics. ^[108] Heavy use of phosphorus fertilizers and their runoff have resulted in eutrophication (overenrichment) of aquatic ecosystems. ^{[109] [110]}

Natural phosphorus-bearing compounds are mostly inaccessible to plants because of the low solubility and mobility in soil. ^[111] Most phosphorus is very stable in the soil minerals or organic matter of the soil. Even when phosphorus is added in manure or fertilizer it can become fixed in the soil. Therefore, the natural phosphorus cycle is very slow. Some of the fixed phosphorus is released again over time, sustaining wild plant growth, however, more is needed to sustain intensive cultivation of crops. ^[112] 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 with calcium phosphate.

Processing phosphate minerals with sulfuric acid for obtaining fertiliser is so important to the global economy that this is the primary industrial market for sulfuric acid and the greatest industrial use of elemental sulfur. ^[113]

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 POCl3 and pesticides
POCl3	Manufacture of plasticiser
P4S10	Manufacturing of additives and pesticides
Na5P3O10	Detergents

Metallurgical aspects

Phosphorus is also an important component in steel production, in the making of phosphor bronze, and in many other related products. ^{[114] [115]} Phosphorus is added to metallic copper during its smelting process to react with oxygen present as an impurity in copper and to produce phosphorus-containing copper (CuOFP) alloys with a higher hydrogen embrittlement resistance than normal copper. ^[116] Phosphate conversion coating is 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 by Charles Sauria in 1830. These matches (and subsequent modifications) were made with heads of white phosphorus, an oxygen-releasing compound (potassium chlorate, lead dioxide, or sometimes nitrate), and a binder. They were poisonous to the workers in manufacture, ^[117] sensitive to storage conditions, toxic if ingested, and hazardous when accidentally ignited on a rough surface. ^{[118] [119]} Production in several countries was banned between 1872 and 1925. ^[120] The international Berne Convention, ratified in 1906, prohibited the use of white phosphorus in matches.

In consequence, phosphorous matches were gradually replaced by safer alternatives. Around 1900 French chemists Henri Sévène and Emile David Cahen invented the modern strike-anywhere match, wherein the white phosphorus was replaced by phosphorus sesquisulfide (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 head potassium chlorate, an oxygen-releasing compound. When struck, small amounts of abrasion from match head and striker strip are mixed intimately to make a small quantity of Armstrong's mixture, a very touch sensitive composition. The fine powder ignites immediately and provides the initial spark to set off the match head. Safety matches separate the two components of the ignition mixture until the match is struck. This is the key safety advantage as it prevents accidental ignition. Nonetheless, safety matches, invented in 1844 by Gustaf Erik Pasch and market ready by the 1860s, did not gain consumer acceptance until the prohibition of white phosphorus. Using a dedicated striker strip was considered clumsy. ^{[48] [104] [121]}

Military

Though military uses of white phosphorus are constrained by modern international law, white phosphorus munitions are still used for military applications, such as incendiary bombs, smoke screens, smoke bombs, and tracer ammunition.

Pesticides and nerve agents

While many organic compounds of phosphorus are required for life, some are highly toxic. A wide range of organophosphorus compounds are used for their toxicity as pesticides and weaponised as nerve agents. ^[31] Some notable examples include sarin, VX or Tabun. Fluorophosphate esters (like sarin) are among the most potent neurotoxins known.

Water softening

Sodium triphosphate is used in laundry detergents in some countries, but banned for this use in others. ^[49] This compound softens the water to enhance the performance of the detergents and to prevent pipe and boiler tube corrosion. ^[122]

Drug production

Phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine. ^[123] For this reason, red and white phosphorus were designated by the United States Drug Enforcement Administration as List I precursor chemicals under 21 CFR 1310.02 effective on November 17, 2001. ^[124] In the United States, handlers of red or white phosphorus are subject to stringent regulatory controls. ^{[125] [126]}

Semiconductors

Phosphorus is a dopant in N-type semiconductors used in high-power electronics and semiconductor detectors. ^[127] In this context, phosphorus is not present at the start of the process, but rather created directly out of silicon during the manufacture of the devices. This is done by neutron transmutation doping, a method based on the conversion of the ³⁰Si into ³¹P by neutron capture and beta decay as follows: $${\displaystyle ^{30}\mathrm {Si} \,(n,\gamma )\,^{31}\mathrm {Si} \rightarrow \,^{31}\mathrm {P} +\beta ^{-}\;(T_{1/2}=2.62\mathrm {h} )}$$

In practice, the silicon is typically placed near or inside a nuclear reactor generating neutrons. As neutrons pass through the silicon, phosphorus atoms are produced by transmutation. This doping method is far less common than diffusion or ion implantation, but it has the advantage of creating an extremely uniform dopant distribution. ^{[128] [129]}

Biological role

See also: Calcium metabolism

Inorganic phosphorus in the form of the phosphate PO3−4 is required for all known forms of life. ^[130] Phosphorus plays a major role in the structural framework of DNA and RNA. Living cells use phosphate to transport cellular energy with adenosine triphosphate (ATP), necessary for every cellular process that uses energy. ATP is also important for phosphorylation, a key regulatory event in cells. Phospholipids are the main structural components of all cellular membranes. Calcium phosphate salts assist in stiffening bones. ^[31] Biochemists commonly use the abbreviation "P_i" to refer to inorganic phosphate. ^[131]

Every living cell is encased in a membrane that separates it from its surroundings. Cellular membranes are composed of a phospholipid matrix and proteins, typically in the form of a bilayer. Phospholipids are derived from glycerol with two of the glycerol hydroxyl (OH) protons replaced by fatty acids as an ester, and the third hydroxyl proton has been replaced with phosphate bonded to another alcohol. ^[132]

An average adult human contains about 0.7 kilograms (1.5 lb) of phosphorus, about 85–90% in bones and teeth in the form of apatite, and the remainder in soft tissues and extracellular fluids. The phosphorus content increases from about 0.5% by mass in infancy to 0.65–1.1% by mass in adults. Average phosphorus concentration in the blood is about 0.4 g/L; about 70% of that is organic and 30% inorganic phosphates. ^[133] 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 as nucleic acids and phospholipids; and excretion almost exclusively in the form of phosphate ions such as H₂PO−4 and HPO2−4. 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 is hydroxyapatite as well as amorphous forms of calcium phosphate, possibly including carbonate. Hydroxyapatite is the main component of tooth enamel. Water fluoridation enhances the resistance of teeth to decay by the partial conversion of this mineral to the still harder material fluorapatite: ^[31]

Ca₅(PO₄)₃OH + F⁻ → Ca₅(PO₄)₃F + OH⁻

Phosphorus deficiency

In medicine, phosphate deficiency syndrome may be caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes that draw phosphate from the blood (such as in refeeding syndrome after malnutrition ^[134] ) or passing too much of it into the urine. All are characterised by hypophosphatemia, which is a condition of low levels of soluble phosphate levels in the blood serum and inside the cells. Symptoms of hypophosphatemia include neurological dysfunction and disruption of muscle and blood cells due to lack of ATP. Too much phosphate can lead to diarrhoea and calcification (hardening) of organs and soft tissue, and can interfere with the body's ability to use iron, calcium, magnesium, and zinc. ^[135]

Phosphorus is an essential macromineral for plants, which is studied extensively in edaphology to understand plant uptake from soil systems. Phosphorus is a limiting factor in many ecosystems; that is, the scarcity of phosphorus limits the rate of organism growth. An excess of phosphorus can also be problematic, especially in aquatic systems where eutrophication sometimes leads to algal blooms. ^[63]

Nutrition

The main food sources for phosphorus are the same as those containing protein, although proteins do not contain phosphorus. For example, milk, meat, and soya typically also have phosphorus. As a rule, if a diet has sufficient protein and calcium, the amount of phosphorus is probably sufficient. ^[136]

The U.S. Institute of Medicine (IOM) updated Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs) for phosphorus in 1997. If there is not sufficient information to establish EARs and RDAs, an estimate designated Adequate Intake (AI) is used instead. The current EAR for phosphorus for people ages 19 and up is 580 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 sets tolerable 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 as Dietary Reference Intakes (DRIs). ^[137]

The European Food Safety Authority (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR. ^[138] AI and UL are defined the same as in the United States. For people ages 15 and older, including pregnancy and lactation, the AI is set at 550 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. ^[139] EFSA reviewed the same safety question and decided that there was not sufficient information to set a UL. ^[140]

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. ^{[141] [142]} A table of the old and new adult daily values is provided at Reference Daily Intake.

Precautions

External contact

NFPA 704 safety square
4 4 2 White phosphorus fire diamond

NFPA 704 safety square
1 1 1 Red phosphorus fire diamond

Elemental phosphorus poses by far the greatest danger in its white form, red phosphorus being relatively nontoxic. ^[143] In the past, external exposure to white phosphorus was treated by washing the affected area with 2% copper(II) sulfate solution to form harmless compounds that are then washed away. According to 2009 United States Navy guidelines: ^[144]

Cupric (copper) sulfate has been used by U.S. personnel in the past and is still being used by some nations. However, copper sulfate is toxic and its use will be discontinued. Copper sulfate may produce kidney and cerebral toxicity as well as intravascular hemolysis.

Instead, the manual suggests:

[...] a bicarbonate solution to neutralise phosphoric acid, which will then allow removal of visible white phosphorus. Particles often can be located by their emission of smoke when air strikes them, or by their phosphorescence in the dark. In dark surroundings, fragments are seen as luminescent spots. Promptly debride the burn if the patient's condition will permit removal of bits of WP (white phosphorus) that might be absorbed later and possibly produce systemic poisoning. DO NOT apply oily-based ointments until it is certain that all WP has been removed. Following complete removal of the particles, treat the lesions as thermal burns.

Ingestion

Because of its common use as a rodenticide, there are documented medical reports of white phosphorus ingestion and its effects, especially on children. ^[145] 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. ^[145]

Passive exposure

Chronic poisoning can lead to necrosis of the jaw. In the United States, exposure to 0.1 mg/m³ of white phosphorus over an 8-hour workday is set as the permissible exposure limit by the Occupational Safety and Health Administration and as the recommended exposure limit by the National Institute for Occupational Safety and Health. From 5 mg/m³, it is considered immediately dangerous to life or health. ^[146]

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