Phosphorus

Phosphorus is a chemical element; it has symbol P and 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 allotropes 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.

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 30.974±0.001 (abridged)
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] 3s^2 3p^3
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)
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/cm^3 red: ≈2.2–2.34 g/cm^3 violet: 2.36 g/cm^3 black: 2.69 g/cm^3
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, −1, 0, +1, +2, +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)
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^−6 cm^3/mol (293 K)
Bulk modulus	white: 5 GPa red: 11 GPa
CAS Number	7723-14-0 (red) 12185-10-3 (white)
History
Naming	from the greek Φωσφόρος, meaning 'light-bearer'
Discovery	Hennig Brand (1669)
Recognised as an element by	Antoine Lavoisier (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

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

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. The discovery is credited to the Hamburg alchemist Hennig Brand in 1669, who was attempting to create the fabled philosopher's stone. To this end, he experimented with urine, which contains considerable quantities of dissolved phosphates from normal metabolism.^: 717 By letting the urine rot (a step later discovered to be unnecessary), 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'. 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.^: 717

Brand at first tried to keep the method secret, but later sold the recipe for 200 thalers to  [de] from Dresden.^: 717 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.^: 717 He was the first to use phosphorus to ignite sulfur-tipped wooden splints, forerunners of modern matches, 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.

In 1777, Antoine Lavoisier recognised phosphorus as an element after Johan Gottlieb Gahn and Carl Wilhelm Scheele showed in 1769 that calcium phosphate is found in bones by obtaining elemental phosphorus from bone ash. Bone ash subsequently became the primary industrial source of phosphorus and remained so until the 1840s. The process consisted of several steps.^: 49–66 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 mining in the Central Chincha Islands, c. 1860

Matches from 1828. The sulfur-tipped match is dipped into liquid containing white phosphorus, and ignites as it is pulled out of the bottle.

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

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

In 1609 Inca Garcilaso de la Vega wrote the book Comentarios Reales in which he described many of the agricultural practices of the Incas prior to the arrival of the Spaniards and introduced the use of guano as a fertiliser. As Garcilaso described, the Incas near the coast harvested guano. 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. 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. 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.

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

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

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 (patented 1889), the use of bone-ash became obsolete.^{: 81–101 : 718–720} After the depletion of world guano sources about the same time, mineral phosphates became the major source of phosphate fertiliser production. Phosphate rock production greatly increased after World War II, and remains the primary global source of phosphorus and phosphorus chemicals today.

The electric furnace method allowed production to increase to the point where it became possible that white phosphorus could be weaponised in 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).^{: 167–185}

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. The difficulty of extinguishing burning phosphorus and the very severe burns it causes had a strong psychological impact on the enemy. Phosphorus incendiary bombs were used on a large scale, notably to destroy Hamburg, the place where the "miraculous bearer of light" was first discovered.

Characteristics

Isotopes

There are 22 known isotopes of phosphorus, ranging from ²⁶P to ⁴⁷P. Only ³¹P is stable and is therefore present at 100% abundance. The half-integer nuclear spin and high abundance of ³¹P make phosphorus-31 nuclear magnetic resonance 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. 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.

Atomic properties

A phosphorus atom has 15 electrons, 5 of which are valence electrons. This results in the electron configuration 1s²2s²2p⁶3s²3p³, often simplified as [Ne]3s²3p³, omitting the core electrons which have a configuration equivalent to the noble gas of the preceding period, in this case neon. The molar ionisation energies of these five electrons are 1011.8, 1907, 2914.1, 4963.6 and 6273.9 kJ⋅mol⁻¹.

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

Allotropes

Crystalline structures of the main phosphorus allotropes

White

Red

Violet

Black

Phosphorus has several allotropes that exhibit very diverse properties. The most useful and therefore common is white phosphorus, followed by red phosphorus. The two other main allotropes, violet and black phosphorus, have either a more fundamental interest or specialised applications. Many other allotropes have been theorised and synthesised, with the search for new materials an active area of research. Commonly mentioned "yellow phosphorus" is not an allotrope, but a result of the gradual degradation of white phosphorus into red phosphorus, accelerated by light and heat. This causes white phosphorus that is aged or otherwise impure (e.g. weapons-grade) to appear yellow.

White phosphorus is a soft, waxy molecular solid that is insoluble in water. It is also very toxic, highly flammable and pyrophoric, igniting in air at about 30 °C (303 K).^{: 721–722} 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). The P₄ molecule in the gas phase has a P-P bond length of 2.1994(3) Å as determined by gas electron diffraction. 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. 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). 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. At still higher temperatures, P₂ dissociates into atomic P.

White phosphorus exposed to air glows in the dark.

When exposed to air, white phosphorus faintly glows green and blue due to oxidation, a phenomenon best visible in the dark. This reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived molecules HPO and P₂O₂ that both emit visible light. However, in a pure-oxygen environment phosphorus does not glow at all, with the oxidation happening only in a range of partial pressures. 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.

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. 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.^: 717 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). After prolonged heating or storage, the color darkens; the resulting product is more stable and does not spontaneously ignite in air.

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" .

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. 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. Single-layer black phosphorus is called phosphorene, and is therefore predictably analogous to graphene.

Natural occurrence

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. In 2020, astronomers analysed ALMA and ROSINA data from the massive star-forming region AFGL 5142, to detect phosphorus-bearing molecules and how they could have been carried in comets to the early Earth.

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

Compounds

Inorganic phosphates

Phoshoric acids

The most prevalent compounds of phosphorus are derivatives of phosphate (PO3−4), 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⁺ + 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)

Food-grade phosphoric acid (additive E338) is used to acidify foods and beverages such as various colas and jams, providing a tangy or sour taste. The phosphoric acid also serves as a preservative. 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, as well as contribute to the formation of kidney stones, especially in those who have had kidney stones previously.

Metal salts

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).

Calcium phosphates in particular are widespread compounds with many applications. Among them, they are used to improve the characteristics of processed meat and cheese, in baking powder, and in toothpaste. Two of the most relevant among them are monocalcium phosphate, and dicalcium phosphate.

Polyphosphates

Phosphate exhibits a tendency to form chains and rings containing P-O-P bonds. Many polyphosphates are known, including 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

Sodium triphosphate 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 and boiler tube corrosion.

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. Although many oxoacids of phosphorus are formed, only nine are commercially important. Among them, hypophosphorous, phosphorous and orthophosphoric 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

Other inorganic compounds

Oxides and sulfides

The tetrahedral structure of P₄O₁₀ and P₄S₁₀

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. Similarly, phosphorus trioxide (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. Mixed oxyhalides and oxyhydrides of phosphorus(III) are almost unknown. Meanwhile, phosphorus forms a wide range of sulfides, where the phosphorus can be in P(V), P(III) or other oxidation states. However, only two of them are commercially significant. Phosphorus pentasulfide (P₄S₁₀) has a structure analogous to P₄O₁₀, and is used in the manufacture of additives and pesticides. The three-fold symmetric Phosphorus sesquisulfide (P₄S₃) is used in strike-anywhere matches.

Halides

Phosphorus halides can have as oxidation state +3 in the case of trihalides and +5 for pentahalides and chalcoalides, but also +2 for disphosphorus tetrahalides. 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 used for the manufacture of pesticides, is produced by chlorination of white phosphorus. The trifluoride is produced from the trichloride by halide exchange. PF₃ is toxic because it binds to haemoglobin.

Most phosphorus pentahalides are common compounds. PF₅ is a colourless gas and the molecules have a trigonal bipyramidal geometry. With fluoride, it forms PF−6, an anion that is isoelectronic with SF₆. PCl₅ is a colourless solid which has an ionic formulation of PCl+4PCl−6, but adopts a trigonal bipyramidal geometry when molten or in the vapour phase. Both the pentafluoride and the pentachloride are Lewis acids. Meanwhile, PBr₅ is an unstable solid formulated as PBr+4Br⁻. PI₅ is not known.

The most important phosphorus oxyhalide is phosphorus oxychloride (POCl₃), which is approximately tetrahedral. It is prepared from PCl₃ and used in the manufacture of plasticizers. Phosphorus can also form thiohalides such as PSCl₃, and in rare cases selenohalides.

Nitrides

The PN molecule phosphorus mononitride is considered unstable, but is a product of crystalline triphosphorus pentanitride decomposition at 1,100 K (830 °C). 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.

Phosphides and phosphine

A wide variety of compounds which contain the containing the phosphide ion P³⁻ exist, both with main-group elements and with metals. They often exhibit complex structures, where phosphorus has the −3 oxidation state. Metal phosphides arise by reaction of metals with red phosphorus. The alkali metals (group 1) and alkaline earth metals (group 2) can also form compounds such as Na₃P₇. These compounds react with water to form phosphine. Some phosphide minerals are also known, like (Fe,Ni)₂P and (Fe,Ni)₃P, but they are very rare on Earth, most instances occurring in iron-nickel meteorites.

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. It is an ill-smelling and toxic gas, 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. The highly flammable gas diphosphine (P₂H₄) is an analogue of hydrazine.

Organophosphorus compounds

A stable diphosphene, a derivative of phosphorus(I)

Phosphines, phosphites and organophosphates

Compounds with P-C and P-O-C bonds are often classified as organophosphorus compounds. They are widely used commercially. The P³⁺ serves as a source of PCl₃ 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

Some organophosphates are used as flame retardants. Among them, 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.

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. Some notable examples include sarin, VX or Tabun. Fluorophosphate esters (like sarin) are among the most potent neurotoxins known.

Thioesters

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, they revert to another phosphorus(III) compound through a sulfonium intermediate.

Phosphorus(I) and phosphorus(II)

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.

Biological role

Cells

Inorganic phosphorus in the form of the phosphate PO3−4 is required for all known forms of life. Phosphorus plays a major role in the structural framework of DNA and RNA. Living cells use phosphate to transport cellular energy with adenosine triphosphate (ATP), necessary for every cellular process that uses energy. ATP is also important for phosphorylation, a key regulatory event in cells. Every living cell is encased in a membrane that separates it from its surroundings. Cellular membranes are composed of a phospholipid matrix and proteins, typically in the form of a bilayer. 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.

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:

Ca₅(PO₄)₃OH + F⁻ → Ca₅(PO₄)₃F + OH⁻

An average adult human contains about 0.7 kilograms (1.5 lb) of phosphorus, about 85–90% in bones and teeth in the form of apatite, and the remainder in soft tissues and extracellular fluids. The phosphorus content increases from about 0.5% by mass in infancy to 0.65–1.1% by mass in adults. In comparison, average phosphorus concentration in the blood is about 0.4 g/L; about 70% of that is organic and 30% inorganic phosphates.

Nutrition

The main food sources for phosphorus are the same as those containing protein, although proteins themselves do not contain phosphorus. For example, milk, meat, and soya typically also have phosphorus. Generally, if a diet includes sufficient protein and calcium, the amount of phosphorus is sufficient.

According to the U.S. Institute of Medicine, the estimated average requirement for phosphorus for people ages 19 and up is 580 mg/day. The RDA is 700 mg/day. RDAs are higher than EARs so as to identify amounts that will cover people with higher-than-average requirements. RDA for pregnancy and lactation are also 700 mg/day. For people ages 1–18 years, the RDA increases with age from 460 to 1250 mg/day. As for safety, the IOM sets tolerable upper intake level for phosphorus at 4000 mg/day. Collectively, these values are referred to as the Dietary Reference Intake. 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. 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. The EFSA reviewed the same safety question and decided that there was not sufficient information to set a UL.

Phosphorus deficiency may be caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes that draw phosphate from the blood (such as in refeeding syndrome after malnutrition) 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.

Phosphorus cycle

Phosphorus is an essential plant nutrient (the most often limiting nutrient, after nitrogen), 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 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 of plant roots, formation of seeds and flowers, and other important factors effecting overall plant health and genetics. Heavy use of phosphorus fertilisers and their runoff have resulted in eutrophication (overenrichment) of aquatic 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 fertiliser it can become fixed in the soil. Therefore, the natural phosphorus cycle is very slow. Some of the fixed phosphorus is released again over time, sustaining wild plant growth, however, more is needed to sustain intensive cultivation of crops. 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.

Production

Mining

Mining of phosphate rock in Nauru

A phosphate train on its way to the port of Casablanca in Morocco.

Means of commercial phosphorus production besides mining are few because the phosphorus cycle does not include significant gas-phase transport. The predominant source of phosphorus in modern times is phosphate rock (as opposed to the guano that preceded it).

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

Processing

Most phosphorus-bearing material is for agriculture fertilisers. In this case where the standards of purity are modest, phosphorus is obtained from phosphate rock by what is called the "wet process." The minerals are treated with sulfuric acid to give phosphoric acid. Phosphoric acid is then neutralised to give various phosphate salts, which comprise fertilisers. In the wet process, phosphorus does not undergo redox. 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.

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. 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:

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.

Reserves

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

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

Global distribution of commercial reserves of rock phosphate in 2016

The countries estimated to have the biggest phosphate rock commercial reserves (in billion metric tons) are Morocco (50), China (3.2), Egypt (2.8), Algeria (2.2), Syria (1.8), Brazil (1.6), Saudi Arabia (1.4), South Africa (1.4), Australia (1.1), United States (1.0), and Finland (1.0). 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. According to some researchers, Earth's commercial and affordable phosphorus reserves are expected to be depleted in 50–100 years.

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. 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. Phosphorus reserve figures are intensely debated. Gilbert suggest that there has been little external verification of the estimate. A 2014 review concluded that the IFDC report "presents an inflated picture of global reserves, in particular those of Morocco, where largely hypothetical and inferred resources have simply been relabeled “reserves".

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 fertiliser application. Strips of grassland or forest between arable land and rivers can also greatly reduce losses of phosphate and other nutrients.

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. 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. However, the technologies currently in use are not yet cost-effective, given the current price of phosphorus on the world market.

Applications

Matches

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

Safety matches are very difficult to ignite on any surface other than a special striker strip. The strip contains non-toxic red phosphorus and the match head potassium chlorate, an oxygen-releasing compound. When struck, small amounts of abrasion 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.

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.

Drug production

Elemental phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine. For this reason, red and white phosphorus are listed in the United States as List I precursor chemicals by the Drug Enforcement Administration, and their handling is subject to stringent regulatory controls.

Metallurgical aspects

Phosphorus is also an important component in steel production, in the making of phosphor 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 higher hydrogen embrittlement resistance than normal copper.Phosphate conversion coating is a chemical treatment applied to steel parts to improve their corrosion resistance.

Semiconductors

Phosphorus is a dopant in N-type semiconductors used in high-power electronics and semiconductor detectors. 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.

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. 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:

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. 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, this very small amount leading to many cases proving fatal, either because of rapid cardiovascular arrest or through the following systemic toxicity.

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.

References

1. "Standard Atomic Weights: Phosphorus". CIAAW. 2013.
2. Prohaska, Thomas; Irrgeher, Johanna; Benefield, Jacqueline; Böhlke, John K.; Chesson, Lesley A.; Coplen, Tyler B.; Ding, Tiping; Dunn, Philip J. H.; Gröning, Manfred; Holden, Norman E.; Meijer, Harro A. J. (2022-05-04). "Standard atomic weights of the elements 2021 (IUPAC Technical Report)". Pure and Applied Chemistry. doi:10.1515/pac-2019-0603. ISSN 1365-3075.
3. Phosphorus at the Encyclopædia Britannica
4. Greenwood, Norman N.; Earnshaw, Alan (1997). Chemistry of the Elements (2nd ed.). Butterworth-Heinemann. p. 28. ISBN 978-0-08-037941-8.
5. Wang, Yuzhong; Xie, Yaoming; Wei, Pingrong; King, R. Bruce; Schaefer, Iii; Schleyer, Paul v. R.; Robinson, Gregory H. (2008). "Carbene-Stabilized Diphosphorus". Journal of the American Chemical Society. 130 (45): 14970–1. Bibcode:2008JAChS.13014970W. doi:10.1021/ja807828t. PMID 18937460.
6. Ellis, Bobby D.; MacDonald, Charles L. B. (2006). "Phosphorus(I) Iodide: A Versatile Metathesis Reagent for the Synthesis of Low Oxidation State Phosphorus Compounds". Inorganic Chemistry. 45 (17): 6864–74. doi:10.1021/ic060186o. PMID 16903744.
7. Arblaster, John W. (2018). Selected Values of the Crystallographic Properties of Elements. Materials Park, Ohio: ASM International. ISBN 978-1-62708-155-9.
8. Lide, D. R., ed. (2005). "Magnetic susceptibility of the elements and inorganic compounds". CRC Handbook of Chemistry and Physics (PDF) (86th ed.). Boca Raton (FL): CRC Press. ISBN 0-8493-0486-5.
9. Weast, Robert (1984). CRC, Handbook of Chemistry and Physics. Boca Raton, Florida: Chemical Rubber Company Publishing. pp. E110. ISBN 0-8493-0464-4.
10. cf. "Memoir on Combustion in General" Mémoires de l'Académie Royale des Sciences 1777, 592–600. from Henry Marshall Leicester and Herbert S. Klickstein, A Source Book in Chemistry 1400–1900 (New York: McGraw Hill, 1952)
11. Weeks, Mary Elvira (1932). "The discovery of the elements. II. Elements known to the alchemists". Journal of Chemical Education. 9 (1): 11. Bibcode:1932JChEd...9...11W. doi:10.1021/ed009p11.
12. Beatty, Richard (2000). Phosphorus. Marshall Cavendish. ISBN 0-7614-0946-7.
13. Parkes, G. D.; Mellor, J. W. (1939). Mellor's Modern Inorganic Chemistry. Longman's Green and Co.
14. Sommers, Michael A. (2007-08-15). Phosphorus. Rosen Group. ISBN 978-1-4042-1960-1.
15. Schmundt, Hilmar (21 April 2010), "Experts Warn of Impending Phosphorus Crisis", Der Spiegel
16. Stillman, J. M. (2003) [1960]. The Story of Alchemy and Early Chemistry. New York: Dover. pp. 418–419. ISBN 0-7661-3230-7.
17. Baccini, Peter; Paul H. Brunner (2012-02-10). Metabolism of the Anthroposphere. MIT Press, 2012. p. 288. ISBN 978-0-262-30054-4.
18. Von Wagner, Rudolf (1897). Manual of chemical technology. New York: D. Appleton & Co. p. 411.
19. Thomson, Robert Dundas (1870). Dictionary of chemistry with its applications to mineralogy, physiology and the arts. Rich. Griffin and Company. p. 416.
20. Threlfall, Richard E. (1951). The Story of 100 years of Phosphorus Making: 1851–1951. Oldbury: Albright & Wilson Ltd.
21. Leigh, G. J. (2004). The World's Greatest Fix: A History of Nitrogen and Agriculture. Oxford University Press. ISBN 978-0-19-516582-1.
22. Skaggs, Jimmy M. (May 1995). The Great Guano Rush: Entrepreneurs and American Overseas Expansion. St. Martin's Press. ISBN 978-0-312-12339-0.
23. Hughes, J. P. W; Baron, R.; Buckland, D. H.; et al. (1962). "Phosphorus Necrosis of the Jaw: A Present-day Study: With Clinical and Biochemical Studies". Br. J. Ind. Med. 19 (2): 83–99. doi:10.1136/oem.19.2.83. PMC 1038164. PMID 14449812.
24. Crass, M. F. Jr. (1941). "A history of the match industry. Part 9" (PDF). Journal of Chemical Education. 18 (9): 428–431. Bibcode:1941JChEd..18..428C. doi:10.1021/ed018p428.^{[permanent dead link]}
25. Oliver, Thomas (1906). "Industrial disease due to certain poisonous fumes or gases". Archives of the Public Health Laboratory. 1. Manchester University Press: 1–21.
26. Charnovitz, Steve (1987). "The Influence of International Labour Standards on the World Trading Regime. A Historical Overview". International Labour Review. 126 (5): 565, 571.
27. Lewis R. Goldfrank; Neal Flomenbaum; Mary Ann Howland; Robert S. Hoffman; Neal A. Lewin; Lewis S. Nelson (2006). Goldfrank's toxicologic emergencies. McGraw-Hill Professional. pp. 1486–1489. ISBN 0-07-143763-0.
28. Toy, Arthur D. F. (1975). The Chemistry of Phosphorus. Texts in Inorganic Chemistry. Vol. 3. Pergamon. ISBN 978-1-4831-4741-3. Retrieved 2013-10-22.
29. US patent 417943
30. Dockery, Kevin (1997). Special Warfare Special Weapons. Chicago: Emperor's Press. ISBN 1-883476-00-3.
31. Greenwood, N. N.; & Earnshaw, A. (1997). Chemistry of the Elements (2nd Edn.), Oxford:Butterworth-Heinemann. ISBN 0-7506-3365-4.
32. Audi, G.; Kondev, F. G.; Wang, M.; Huang, W. J.; Naimi, S. (2017). "The NUBASE2016 evaluation of nuclear properties" (PDF). Chinese Physics C. 41 (3): 030001. Bibcode:2017ChPhC..41c0001A. doi:10.1088/1674-1137/41/3/030001.
33. Neufcourt, L.; Cao, Y.; Nazarewicz, W.; Olsen, E.; Viens, F. (2019). "Neutron drip line in the Ca region from Bayesian model averaging". Physical Review Letters. 122 (6): 062502–1–062502–6. arXiv:1901.07632. Bibcode:2019PhRvL.122f2502N. doi:10.1103/PhysRevLett.122.062502. PMID 30822058. S2CID 73508148.
34. David A. Atwood, ed. (2013-02-19). Radionuclides in the Environment. John Wiley & Sons, 2013. ISBN 978-1-118-63269-7.
35. "Phosphorus-32" (PDF). University of Michigan Department of Occupational Safety & Environmental Health. Archived from the original (PDF) on 2016-05-28. Retrieved 2010-11-18.
36. A. Holleman; N. Wiberg (1985). "XV 2.1.3". Lehrbuch der Anorganischen Chemie (33rd ed.). de Gruyter. ISBN 3-11-012641-9.
37. Tian, Haijiang; Wang, Jiahong; Lai, Gengchang; Dou, Yanpeng; Gao, Jie; Duan, Zunbin; Feng, Xiaoxiao; Wu, Qi; He, Xingchen; Yao, Linlin; Zeng, Li; Liu, Yanna; Yang, Xiaoxi; Zhao, Jing; Zhuang, Shulin; Shi, Jianbo; Qu, Guangbo; Yu, Xue-Feng; Chu, Paul K.; Jiang, Guibin (2023). "Renaissance of elemental phosphorus materials: properties, synthesis, and applications in sustainable energy and environment". Chemical Society Reviews. 52 (16): 5388–5484. doi:10.1039/D2CS01018F. ISSN 0306-0012. PMID 37455613. Retrieved 2025-02-25.
38. Cossairt, Brandi M.; Cummins, Christopher C.; Head, Ashley R.; Lichtenberger, Dennis L.; Berger, Raphael J. F.; Hayes, Stuart A.; Mitzel, Norbert W.; Wu, Gang (2010-06-01). "On the Molecular and Electronic Structures of AsP3 and P4". Journal of the American Chemical Society. 132 (24): 8459–8465. Bibcode:2010JAChS.132.8459C. doi:10.1021/ja102580d. ISSN 0002-7863. PMID 20515032.
39. Welford C. Roberts; William R. Hartley (1992-06-16). Drinking Water Health Advisory: Munitions (illustrated ed.). CRC Press, 1992. p. 399. ISBN 0-87371-754-6.
40. Marie-Thérèse Averbuch-Pouchot; A. Durif (1996). Topics in Phosphate Chemistry. World Scientific, 1996. p. 3. ISBN 981-02-2634-9.
41. Simon, Arndt; Borrmann, Horst; Horakh, Jörg (1997). "On the Polymorphism of White Phosphorus". Chemische Berichte. 130 (9): 1235–1240. doi:10.1002/cber.19971300911.
42. Piro, N. A.; Figueroa, J. S.; McKellar, J. T.; Cummins, C. C. (2006). "Triple-Bond Reactivity of Diphosphorus Molecules". Science. 313 (5791): 1276–9. Bibcode:2006Sci...313.1276P. doi:10.1126/science.1129630. PMID 16946068. S2CID 27740669.
43. Vanzee, Richard J.; Khan, Ahsan U. (1976). "The phosphorescence of phosphorus". The Journal of Physical Chemistry. 80 (20): 2240–2242. doi:10.1021/j100561a021.
44. "Nobel Prize in Chemistry 1956 – Presentation Speech by Professor A. Ölander (committee member)". Retrieved 2009-05-05.
45. Shen, Z; Yu, JC (2016). "Nanostructured elemental photocatalysts: Development and challenges". In Yamashita, H; Li, H (eds.). Nanostructured Photocatalysts: Advanced Functional Materials. Switzerland: Springer. pp. 295–312 (301). ISBN 978-3-319-26077-8.
46. Egon Wiberg; Nils Wiberg; Arnold Frederick Holleman (2001). Inorganic chemistry. Academic Press. pp. 683–684, 689. ISBN 978-0-12-352651-9. Retrieved 2011-11-19.
47. Hammond, C. R. (2000). "The Elements". Handbook of Chemistry and Physics (81st ed.). CRC press. ISBN 0-8493-0481-4.
48. Berger, L. I. (1996). Semiconductor materials. CRC Press. p. 84. ISBN 0-8493-8912-7.
49. Robert Engel (2003-12-18). Synthesis of Carbon-Phosphorus Bonds (2 ed.). CRC Press, 2003. p. 11. ISBN 0-203-99824-3.
50. A. Brown; S. Runquist (1965). "Refinement of the crystal structure of black phosphorus". Acta Crystallogr. 19 (4): 684–685. Bibcode:1965AcCry..19..684B. doi:10.1107/S0365110X65004140.
51. Cartz, L.; Srinivasa, S.R.; Riedner, R.J.; Jorgensen, J.D.; Worlton, T.G. (1979). "Effect of pressure on bonding in black phosphorus". Journal of Chemical Physics. 71 (4): 1718–1721. Bibcode:1979JChPh..71.1718C. doi:10.1063/1.438523.
52. Lange, Stefan; Schmidt, Peer & Nilges, Tom (2007). "Au3SnP7@Black Phosphorus: An Easy Access to Black Phosphorus". Inorg. Chem. 46 (10): 4028–35. doi:10.1021/ic062192q. PMID 17439206.
53. Koo, B.-C.; Lee, Y.-H.; Moon, D.-S.; Yoon, S.-C.; Raymond, J. C. (2013). "Phosphorus in the Young Supernova Remnant Cassiopeia A". Science. 342 (6164): 1346–8. arXiv:1312.3807. Bibcode:2013Sci...342.1346K. doi:10.1126/science.1243823. PMID 24337291. S2CID 35593706.
54. Rivilla, V. M.; Drozdovskaya, M. N.; Altwegg, K.; Caselli, P.; Beltrán, M. T.; Fontani, F.; van der Tak, F. F. S.; Cesaroni, R.; Vasyunin, A.; Rubin, M.; Lique, F.; Marinakis, S.; Testi, L. (2019). "ALMA and ROSINA detections of phosphorus-bearing molecules: the interstellar thread between star-forming regions and comets". Monthly Notices of the Royal Astronomical Society. 492: 1180–1198. arXiv:1911.11647. doi:10.1093/mnras/stz3336. S2CID 208290964.
55. ‹See TfM›reference=Corbridge, D.E.C. (1995) "Phosphorus: An Outline of its Chemistry, Biochemistry, and Technology" 5th Edition Elsevier: Amsterdam. ISBN 0-444-89307-5
56. "Current EU approved additives and their E Numbers". Foods Standards Agency. 14 March 2012. Archived from the original on 21 August 2013. Retrieved 22 July 2012.
57. "Why is phosphoric acid used in some Coca‑Cola drinks?| Frequently Asked Questions | Coca-Cola GB". www.coca-cola.co.uk. Archived from the original on 2 August 2021. Retrieved 2021-08-31.
58. Moynihan, P. J. (23 November 2002). "Dietary advice in dental practice". British Dental Journal. 193 (10): 563–568. doi:10.1038/sj.bdj.4801628. PMID 12481178.
59. Qaseem, A; Dallas, P; Forciea, MA; Starkey, M; et al. (4 November 2014). "Dietary and pharmacologic management to prevent recurrent nephrolithiasis in adults: A clinical practice guideline from the American College of Physicians". Annals of Internal Medicine. 161 (9): 659–67. doi:10.7326/M13-2908. PMID 25364887.
60. 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. ISBN 978-3-527-30673-2.
61. Heal, H. G. "The Inorganic Heterocyclic Chemistry of Sulfur, Nitrogen, and Phosphorus" Academic Press: London; 1980. ISBN 0-12-335680-6.
62. Mark, J. E.; Allcock, H. R.; West, R. "Inorganic Polymers" Prentice Hall, Englewood, NJ: 1992. ISBN 0-13-465881-7.
63. Naiker, Vidhukrishnan E.; Mestry, Siddhesh; Nirgude, Tejal; Gadgeel, Arjit; Mhaske, S. T. (2023-01-01). "Recent developments in phosphorous-containing bio-based flame-retardant (FR) materials for coatings: an attentive review". Journal of Coatings Technology and Research. 20 (1): 113–139. doi:10.1007/s11998-022-00685-z. ISSN 1935-3804. S2CID 253349703.
64. Diskowski, Herbert; Hofmann, Thomas. "Phosphorus". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a19_505. ISBN 978-3-527-30673-2.
65. Almasi, Lucreţia (1971). "The Sulfur–Phosphorus Bond". In Senning, Alexander (ed.). Sulfur in Organic and Inorganic Chemistry. Vol. 1. New York: Marcel Dekker. pp. 49–55. ISBN 0-8247-1615-9. LCCN 70-154612.
66. ‹See TfM›reference=Ruttenberg, K. C. Phosphorus Cycle – Terrestrial Phosphorus Cycle, Transport of Phosphorus, from Continents to the Ocean, The Marine Phosphorus Cycle. Archived July 13, 2011, at archive.today
67. Nelson, D. L.; Cox, M. M. "Lehninger, Principles of Biochemistry" 3rd Ed. Worth Publishing: New York, 2000. ISBN 1-57259-153-6.
68. Bernhardt, Nancy E.; Kasko, Artur M. (2008). Nutrition for the Middle Aged and Elderly. Nova Publishers. p. 171. ISBN 978-1-60456-146-3.
69. ‹See TfM›reference=Phosphorus in diet: MedlinePlus Medical Encyclopedia. Nlm.nih.gov (2011-11-07). Retrieved on 2011-11-19.
70. Institute of Medicine (1997). "Phosphorus". Dietary Reference Intakes for Calcium, Phosphorus, Magnesium, Vitamin D, and Fluoride. Washington, DC: The National Academies Press. pp. 146–189. doi:10.17226/5776. ISBN 978-0-309-06403-3. PMID 23115811. S2CID 8768378.
71. "Dietary reference values". European Food Safety Authority. 7 August 2024. Retrieved 24 September 2024.
72. "Overview on Dietary Reference Values for the EU population as derived by the EFSA Panel on Dietetic Products, Nutrition and Allergies" (PDF). 2017.
73. "Tolerable Upper Intake Levels For Vitamins And Minerals" (PDF). European Food Safety Authority. 2006. Retrieved 2025-03-22.
74. Mehanna H. M.; Moledina J.; Travis J. (June 2008). "Refeeding syndrome: what it is, and how to prevent and treat it". BMJ. 336 (7659): 1495–8. doi:10.1136/bmj.a301. PMC 2440847. PMID 18583681.
75. Anderson, John J. B. (1996). "Calcium, Phosphorus and Human Bone Development". Journal of Nutrition. 126 (4 Suppl): 1153S – 1158S. doi:10.1093/jn/126.suppl_4.1153S. PMID 8642449.
76. Etesami, H. (2019). Nutrient Dynamics for Sustainable Crop Production. Springer. p. 217. ISBN 978-981-13-8660-2.
77. Philpott, Tom (March–April 2013). "You Need Phosphorus to Live—and We're Running Out". Mother Jones.
78. Carpenter, Stephen R. (2005). "Eutrophication of aquatic ecosystems: Bistability and soil phosphorus". Proceedings of the National Academy of Sciences. 102 (29): 10002–10005. Bibcode:2005PNAS..10210002C. doi:10.1073/pnas.0503959102. ISSN 0027-8424. PMC 1177388. PMID 15972805.
79. Conley, Daniel J.; Paerl, Hans W.; Howarth, Robert W.; et al. (2009). "Controlling Eutrophication: Nitrogen and Phosphorus". Science. 323 (5917): 1014–1015. doi:10.1126/science.1167755. ISSN 0036-8075. PMID 19229022.
80. "Soil Phosphorous" (PDF). United States Department of Agriculture. Archived from the original (PDF) on 2020-10-28. Retrieved 2020-08-17.
81. "Managing Phosphorus for Crop Production". Penn State Extension. Archived from the original on 2020-10-20. Retrieved 2020-08-17.
82. Jessica Elzea Kogel, ed. (2006). Industrial Minerals & Rocks: Commodities, Markets, and Uses. SME, 2006. p. 964. ISBN 0-87335-233-5.
83. Neset, Tina-Simone S.; Cordell, Dana (2011). "Global phosphorus scarcity: identifying synergies for a sustainable future". Journal of the Science of Food and Agriculture. 92 (1): 2–6. doi:10.1002/jsfa.4650. PMID 21969145.
84. US Geological Survey, Phosphate Rock, 2021.
85. Geeson, Michael B.; Cummins, Christopher C. (2020). "Let's Make White Phosphorus Obsolete". ACS Central Science. 6 (6): 848–860. doi:10.1021/acscentsci.0c00332. PMC 7318074. PMID 32607432.
86. Tayibi, Hanan; Choura, Mohamed; López, Félix A.; Alguacil, Francisco J.; López-Delgado, Aurora (2009). "Environmental Impact and Management of Phosphogypsum". Journal of Environmental Management. 90 (8): 2377–2386. Bibcode:2009JEnvM..90.2377T. doi:10.1016/j.jenvman.2009.03.007. hdl:10261/45241. PMID 19406560.
87. Shriver, Atkins. Inorganic Chemistry, Fifth Edition. W. H. Freeman and Company, New York; 2010; p. 379.
88. "ERCO and Long Harbour". Memorial University of Newfoundland and the C.R.B. Foundation. Retrieved 2009-06-06.
89. "Phosphate Rock Statistics and Information U.S. Geological Survey". www.usgs.gov. Retrieved 2023-04-09.
90. American Geophysical Union, Fall Meeting 2007, abstract #V33A-1161. Mass and Composition of the Continental Crust
91. Sutton, M.A.; Bleeker, A.; Howard, C.M.; et al. (2013). Our Nutrient World: The challenge to produce more food and energy with less pollution (PDF). Centre for Ecology and Hydrology, Edinburgh on behalf of the Global Partnership on Nutrient Management and the International Nitrogen Initiative. ISBN 978-1-906698-40-9. Archived from the original (PDF) on 2016-11-04. Retrieved 2015-05-12.
92. CIM DEFINITION STANDARDS - For Mineral Resources and Mineral Reserves (PDF). CIM Standing Committee on Reserve Definitions. 2010. pp. 4–6. Archived from the original (PDF) on 14 February 2019. Retrieved 5 April 2024.
93. Amundson, R.; Berhe, A. A.; Hopmans, J. W.; Olson, C.; Sztein, A. E.; Sparks, D. L. (2015). "Soil and human security in the 21st century". Science. 348 (6235): 1261071. Bibcode:2015Sci...34861071A. doi:10.1126/science.1261071. ISSN 0036-8075. PMID 25954014. S2CID 206562728.
94. Ahokas, K. (2015). "Finland's phosphorus resources are more important than ever (Geological Survey of Finland)". Archived from the original on 2019-05-06. Retrieved 2017-04-01.
95. "Phosphate rock, U.S. Geological Survey, Mineral Commodity Summaries, January 2025" (PDF). Retrieved 2025-03-05.
96. Walan, P.; Davidsson, S.; Johansson, S.; Höök, M. (2014). "Phosphate rock production and depletion: Regional disaggregated modeling and global implications". Resources, Conservation and Recycling. 93 (12): 178–187. Bibcode:2014RCR....93..178W. doi:10.1016/j.resconrec.2014.10.011. Retrieved 9 October 2017.
97. Cordell, Dana; Drangert, Jan-Olof; White, Stuart (2009). "The story of phosphorus: Global food security and food for thought". Global Environmental Change. 19 (2): 292–305. Bibcode:2009GEC....19..292C. doi:10.1016/j.gloenvcha.2008.10.009. ISSN 0959-3780. S2CID 1450932.
98. "IFDC Report Indicates Adequate Phosphorus Resources Available to Meet Global Food Demands". 22 September 2010.
99. Van Kauwenbergh, Steven J. (2010). World Phosphate Rock Reserves and Resources. Muscle Shoals, AL, USA: International Fertilizer Development Center (IFDC). p. 60. ISBN 978-0-88090-167-3. Archived from the original on 19 August 2018. Retrieved 7 April 2016.
100. Van Vuuren, D.P.; Bouwman, A.F.; Beusen, A.H.W. (2010). "Phosphorus demand for the 1970–2100 period: A scenario analysis of resource depletion". Global Environmental Change. 20 (3): 428–439. Bibcode:2010GEC....20..428V. doi:10.1016/j.gloenvcha.2010.04.004. ISSN 0959-3780.
101. Gilbert, Natasha (8 October 2009). "The disappearing nutrient". Nature. 461 (7265): 716–718. doi:10.1038/461716a. PMID 19812648. S2CID 4419892.
102. Edixhoven, J. D.; Gupta, J.; Savenije, H. H. G. (2014). "Recent revisions of phosphate rock reserves and resources: a critique" (PDF). Earth System Dynamics. 5 (2): 491–507. Bibcode:2014ESD.....5..491E. doi:10.5194/esd-5-491-2014. ISSN 2190-4987.
103. Udawatta, Ranjith P.; Henderson, Gray S.; Jones, John R.; Hammer, David (2011). "Phosphorus and nitrogen losses in relation to forest, pasture and row-crop land use and precipitation distribution in the midwest usa". Journal of Water Science. 24 (3): 269–281. doi:10.7202/1006477ar.
104. Tweed, Katherine (2009-11-01). "Sewage Industry Fights Phosphorus Pollution". Scientific American. Retrieved 2024-06-21.
105. Sartorius, C., von Horn, J., Tettenborn, F. (2011). Phosphorus recovery from wastewater – state-of-the-art and future potential. Conference presentation at Nutrient Recovery and Management Conference organised by International Water Association (IWA) and Water Environment Federation (WEF) in Florida, USA
106. Alexander P. Hardt (2001). "Matches". Pyrotechnics. Post Falls Idaho US: Pyrotechnica Publications. pp. 74–84. ISBN 0-929388-06-2.
107. Skinner, H.F. (1990). "Methamphetamine synthesis via hydriodic acid/red phosphorus reduction of ephedrine". Forensic Science International. 48 (2): 123–134. doi:10.1016/0379-0738(90)90104-7.
108. "66 FR 52670—52675". 17 October 2001. Retrieved 2009-05-05.
109. "21 cfr 1309". Archived from the original on 2009-05-03. Retrieved 2009-05-05.
110. "21 USC, Chapter 13 (Controlled Substances Act)". Retrieved 2009-05-05.
111. Roland W. Scholz; Amit H. Roy; Fridolin S. Brand; Deborah Hellums; Andrea E. Ulrich, eds. (2014-03-12). Sustainable Phosphorus Management: A Global Transdisciplinary Roadmap. Springer Science & Business Media. p. 175. ISBN 978-94-007-7250-2.
112. Mel Schwartz (2016-07-06). Encyclopedia and Handbook of Materials, Parts and Finishes. CRC Press. ISBN 978-1-138-03206-4.
113. Joseph R. Davisz, ed. (January 2001). Copper and Copper Alloys. ASM International. p. 181. ISBN 0-87170-726-8.
114. "Neutron Transmutation Doping of Silicon | MIT Nuclear Reactor Laboratory".
115. Baliga, B. Jayant (1987-03-10). Modern Power Devices. Wiley-Interscience. p. 32. ISBN 0-471-81986-7.
116. Schmidt, P. E.; Vedde, J. (1998). High Resistivity NTD Production and Applications. Electrochemical Society Proceedings. Vol. 98. ISBN 9781566772075.
117. "Public Health England" (PDF). Public Health England. January 2016. Retrieved 2 March 2025.
118. "US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries". Archived from the original on November 22, 2005. Retrieved 2009-05-05.
119. Simon, Frank A. (1976-03-29). "Acute Yellow Phosphorus Poisoning: "Smoking Stool Syndrome"". JAMA. 235 (13): 1343–1344. doi:10.1001/jama.1976.03260390029021. ISSN 0098-7484. PMID 946251.
120. "CDC - NIOSH Pocket Guide to Chemical Hazards - Phosphorus (yellow)". www.cdc.gov. Retrieved 2015-11-21.