Phosphorus

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
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
Discovery	Hennig Brand (1669)
Recognised as an element by	Antoine Lavoisier[9] (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

.

Elemental phosphorus was first isolated as white phosphorus in 1669. In white phosphorus, phosphorus atoms are arranged in groups of 4, written as P₄. White phosphorus emits a faint glow when exposed to

.

Phosphorus is an element essential to sustaining

.

Characteristics

Allotropes

Phosphorus has several allotropes that exhibit strikingly diverse properties.^[10] The two most common allotropes are white phosphorus and red phosphorus.^[11]

For both pure and applied uses, the most important allotrope

White phosphorus exists in two crystalline forms: α (alpha) and β (beta). At room temperature, the α-form is stable. It is more common, has cubic crystal structure and at 195.2 K (−78.0 °C), it transforms into β-form, which has hexagonal crystal structure. These forms differ in terms of the relative orientations of the constituent P₄ tetrahedra.^[14][15]

White phosphorus is the least stable, the most reactive, the most

Thermal decomposition of P₄ at 1100 K gives diphosphorus, P₂. This species is not stable as a solid or liquid. The dimeric unit contains a triple bond and is analogous to N₂. It can also be generated as a transient intermediate in solution by thermolysis of organophosphorus precursor reagents.^[17] At still higher temperatures, P₂ dissociates into atomic P.^[16]

Properties of some allotropes of phosphorus^[10][18]

Form	white(α)	white(β)	red	violet	black
Symmetry	Body-centred cubic	Triclinic	Amorphous	Monoclinic	Orthorhombic
Pearson symbol		aP24		mP84	oS8
Space group	I43m	P1 No.2		P2/c No.13	Cmce No.64
Density (g/cm3)	1.828	1.88	~2.2	2.36	2.69
Band gap (eV)	2.1		1.8	1.5	0.34
Refractive index	1.8244			2.6	2.4

Another form, scarlet phosphorus, is obtained by allowing a solution of white phosphorus in carbon disulfide to evaporate in sunlight.^[18]

Chemiluminescence

When first isolated, it was observed that the green glow emanating from white phosphorus would persist for a time in a stoppered jar, but then cease. Robert Boyle in the 1680s ascribed it to "debilitation" of the air. Actually, it is oxygen being consumed. By the 18th century, it was known that in pure oxygen, phosphorus does not glow at all;^[28] there is only a range of partial pressures at which it does. Heat can be applied to drive the reaction at higher pressures.^[29]

In 1974, the glow was explained by R. J. van Zee and A. U. Khan.^[30][31] A reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming the short-lived molecules HPO and P
₂O
₂ that both emit visible light. The reaction is slow and only very little of the intermediates are required to produce the luminescence, hence the extended time the glow continues in a stoppered jar.

Since its discovery, phosphors and phosphorescence were used loosely to describe substances that shine in the dark without burning. Although the term phosphorescence is derived from phosphorus, the reaction that gives phosphorus its glow is properly called chemiluminescence (glowing due to a cold chemical reaction), not phosphorescence (re-emitting light that previously fell onto a substance and excited it).^[32]

Isotopes

There are 22 known

spectroscopy a very useful analytical tool in studies of phosphorus-containing samples.

Two

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

• radiolabeled DNA and RNA probes, e.g. for use in Northern blots or Southern blots
  .
• ³³
  P, a beta-emitter (0.25 MeV) with a half-life of 25.4 days. It is used in life-science laboratories in applications in which lower energy beta emissions are advantageous such as DNA sequencing.

The high-energy beta particles from ³²
P penetrate skin and corneas and any ³²
P ingested, inhaled, or absorbed is readily incorporated into bone and nucleic acids. For these reasons, Occupational Safety and Health Administration in the United States, and similar institutions in other developed countries require personnel working with ³²
P to wear lab coats, disposable gloves, and safety glasses or goggles to protect the eyes, and avoid working directly over open containers. Monitoring personal, clothing, and surface contamination is also required. Shielding requires special consideration. 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 acrylic or other plastic, water, or (when transparency is not required), even wood.^[35]

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.^[36]

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.^[37][38]

Crust and organic sources

Phosphorus has a concentration in the Earth's crust of about one gram per kilogram (compare copper at about 0.06 grams). It is not found free in nature, but is widely distributed in many

Organic sources, namely

.

Compounds

Phosphorus(V)

The most prevalent compounds of phosphorus are derivatives of phosphate (PO₄³⁻), a tetrahedral anion.^[44] Phosphate is the conjugate base of phosphoric acid, which is produced on a massive scale for use in fertilisers. Being triprotic, phosphoric acid converts stepwise to three conjugate bases:

H₃PO₄ + H₂O ⇌ H₃O⁺ + H₂PO₄⁻       K_a1 = 7.25×10⁻³

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

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

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

:

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

, phosphate forms a variety of salts. These solids are polymeric, featuring P-O-M linkages. When the metal cation has a charge of 2+ or 3+, the salts are generally insoluble, hence they exist as common minerals. Many phosphate salts are derived from hydrogen phosphate (HPO₄²⁻).

, (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.^[45]

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

.

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.

Phosphorus(I) and phosphorus(II)

These compounds generally feature P–P bonds.^[16] 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, P³⁻. 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.^[16] Schreibersite is a naturally occurring metal-rich phosphide found in meteorites. The structures of the metal-rich and phosphorus-rich phosphides can be complex.

.

Oxoacids

Phosphorus

, 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.^[46]

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.^[47] Mixed oxyhalides and oxyhydrides of phosphorus(III) are almost unknown.

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

:

PCl₃ + 3 C₆H₅OH → P(OC₆H₅)₃ + 3 HCl

Similar reactions occur for

:

OPCl₃ + 3 C₆H₅OH → OP(OC₆H₅)₃ + 3 HCl

History

Etymology

The name Phosphorus in Ancient Greece was the name for the planet Venus and is derived from the Greek words (φῶς = light, φέρω = carry), which roughly translates as light-bringer or light carrier.^[20] (In Greek mythology and tradition, Augerinus (Αυγερινός = morning star, still in use today), Hesperus or Hesperinus (΄Εσπερος or Εσπερινός or Αποσπερίτης = evening star, still in use today) and Eosphorus (Εωσφόρος = dawnbearer, not in use for the planet after Christianity) are close homologues, and also associated with Phosphorus-the-morning-star).

According to the Oxford English Dictionary, the correct spelling of the element is phosphorus. The word phosphorous is the adjectival form of the P³⁺ valence: so, just as sulfur forms sulfurous and sulfuric compounds, phosphorus forms phosphorous compounds (e.g., phosphorous acid) and P⁵⁺ valence phosphoric compounds (e.g., phosphoric acids and phosphates).

Discovery

The discovery of phosphorus, the first element to be discovered that was not known since ancient times,^[48] is credited to the German alchemist Hennig Brand in 1669, although others might have discovered phosphorus around the same time.^[49] Brand experimented with urine, which contains considerable quantities of dissolved phosphates from normal metabolism.^[20] Working in Hamburg, Brand attempted to create the fabled philosopher's stone through the distillation of some salts by evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. It was named phosphorus mirabilis ("miraculous bearer of light").^[50]

Brand's process originally involved letting urine stand for days until it gave off a terrible smell. Then he boiled it down to a paste, heated this paste to a high temperature, and led the vapours through water, where he hoped they would condense to gold. Instead, he obtained a white, waxy substance that glowed in the dark. Brand had discovered phosphorus. Specifically, Brand produced ammonium sodium hydrogen phosphate, (NH
₄)NaHPO
₄. While the quantities were essentially correct (it took about 1,100 litres [290 US gal] of urine to make about 60 g of phosphorus), it was unnecessary to allow the urine to rot first. Later scientists discovered that fresh urine yielded the same amount of phosphorus.^[32]

Brand at first tried to keep the method secret,^[51] but later sold the recipe for 200 thalers to Johann Daniel Kraft (de) from Dresden.^[20] Krafft toured much of Europe with it, including England, where he met with Robert Boyle. The secret—that the substance was made from urine—leaked out, and Johann Kunckel (1630–1703) was able to reproduce it in Sweden (1678). Later, Boyle in London (1680) also managed to make phosphorus, possibly with the aid of his assistant, Ambrose Godfrey-Hanckwitz. Godfrey later made a business of the manufacture of phosphorus.

Boyle states that Krafft gave him no information as to the preparation of phosphorus other than that it was derived from "somewhat that belonged to the body of man". This gave Boyle a valuable clue, so that he, too, managed to make phosphorus, and published the method of its manufacture.^[20] Later he improved Brand's process by using sand in the reaction (still using urine as base material),

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

Robert Boyle was the first to use phosphorus to ignite sulfur-tipped wooden splints, forerunners of our modern matches, in 1680.^[52]

Phosphorus was the 13th element to be discovered. Because of its tendency to spontaneously combust when left alone in air, it is sometimes referred to as "the Devil's element".^[53]

Bone ash and guano

Antoine Lavoisier recognized phosphorus as an element in 1777 after Johan Gottlieb Gahn and Carl Wilhelm Scheele, in 1769, showed that calcium phosphate (Ca
₃(PO
₄)
₂) is found in bones by obtaining elemental phosphorus from bone ash.^[9]

Bone ash was the major source of phosphorus until the 1840s. The method started by roasting bones, then employed the use of

.

In the 1840s, world phosphate production turned to the mining of tropical island deposits formed from bird and bat guano (see also Guano Islands Act). These became an important source of phosphates for fertiliser in the latter half of the 19th century.^[56]

Phosphate rock

" fertiliser products.

Incendiaries

White phosphorus was first made commercially in the 19th century for the

Early matches used white phosphorus in their composition, which was dangerous due to its toxicity. Murders, suicides and accidental

Production

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.^[64] Critical to contemporary agriculture, its annual demand is rising nearly twice as fast as the growth of the human population.^[40] The production of phosphorus may have peaked before 2011 and some scientists predict reserves will be depleted before the end of the 21st century.^[65][40][66] Phosphorus comprises about 0.1% by mass of the average rock, and consequently, the Earth's supply is vast, though dilute.^[16]

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.^[67] 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.^[68]

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.^[67] 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 SiO
₂, and coke to produce P
₄. The P
₄ product, being volatile, is readily isolated:^[69]

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.^[70]

Historical routes

Historically, before the development of mineral-based extractions, white phosphorus was isolated on an industrial scale from bone ash.^[71] In this process, the tricalcium phosphate in bone ash is converted to monocalcium phosphate with sulfuric acid:

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

Monocalcium phosphate is then dehydrated to the corresponding metaphosphate:

Ca(H₂PO₄)₂ → Ca(PO₃)₂ + 2 H₂O

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

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

Peak phosphorus

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

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][65] 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

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

The accurate determination of peak phosphorus is dependent on knowing the total world's commercial

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

In 2023, the United States Geological Survey (USGS) estimated that economically extractable phosphate rock reserves worldwide are 72 billion tons, while world mining production in 2022 was 220 million tons.^[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][89][90] Gilbert suggest that there has been little external verification of the estimate.^[91] 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.^[92][77]

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.^[93]

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.^[94]

United States

US production of phosphate rock peaked in 1980 at 54.4 million metric tons. The United States was the world's largest producer of phosphate rock from at least 1900, up until 2006, when US production was exceeded by that of China. In 2019, the US produced 10 percent of the world's phosphate rock.^[95]

Exhaustion of guano reserves

In 1609

.

Phosphorus conservation and recycling

Overview

Phosphorus can be transferred from the soil in one location to another as food is transported across the world, taking the phosphorus it contains with it. Once consumed by humans, it can end up in the local environment (in the case of

.

In an effort to postpone the onset of peak phosphorus several methods of reducing and reusing phosphorus are in practice, such as in agriculture and in sanitation systems. The Soil Association, the UK organic agriculture certification and pressure group, issued a report in 2010 "A Rock and a Hard Place" encouraging more recycling of phosphorus.^[98] One potential solution to the shortage of phosphorus is greater recycling of human and animal wastes back into the environment.^[99]

Agricultural practices

Reducing agricultural runoff and soil erosion can slow the frequency with which farmers have to reapply phosphorus to their fields. Agricultural methods such as

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

Excreta reuse

Main article:

Excreta reuse

The oldest method of recycling phosphorus is through the reuse of animal

excreta

in agriculture. Via this method, phosphorus in the foods consumed are excreted, and the animal or human excreta are subsequently collected and re-applied to the fields. Although this method has maintained civilizations for centuries the current system of manure management is not logistically geared towards application to crop fields on a large scale. At present, manure application could not meet the phosphorus needs of large scale agriculture. Despite that, it is still an efficient method of recycling used phosphorus and returning it to the soil.

Sewage sludge

Sewage treatment plants that have an enhanced biological phosphorus removal step produce a sewage sludge that is rich in phosphorus. Various processes have been developed to extract phosphorus from sewage sludge directly, from the ash after incineration of the sewage sludge or from other products of sewage sludge treatment. This includes the extraction of phosphorus rich materials such as struvite from waste processing plants.^[91] The struvite can be made by adding magnesium to the waste. Some companies such as Ostara in Canada and NuReSys in Belgium are already using this technique to recover phosphate. Ostara has eight operating plants worldwide.^{[citation needed]}

Research on phosphorus recovery methods from sewage sludge has been carried out in Sweden and Germany since around 2003, but the technologies currently under development are not yet cost effective, given the current price of phosphorus on the world market.^[101][102]

Applications

Flame retardant

Phosphorus compounds are used as flame retardants. Flame-retardant materials and coatings are being developed that are both phosphorus- and bio-based.^[103]

Food additive

Phosphorus is an essential mineral for humans listed in the Dietary Reference Intake (DRI).

Food-grade phosphoric acid (additive E338^[104]) 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.^[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] Phosphoric acid also has the potential to contribute to the formation of kidney stones, especially in those who have had kidney stones previously.^[107]

Fertiliser

Phosphorus is an essential plant nutrient (the most often limiting nutrient, after

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

Organophosphorus

White phosphorus is widely used to make

Metallurgical aspects

Phosphorus is also an important component in

Phosphate conversion coating is a chemical treatment applied to steel parts to improve their corrosion resistance.

Matches

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,^[119] sensitive to storage conditions, toxic if ingested, and hazardous when accidentally ignited on a rough surface.^[120][121] Production in several countries was banned between 1872 and 1925.^[122] 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.^{[21][114][123]}

Water softening

Miscellaneous

• Phosphates are used to make special glasses for
  sodium lamps.^[23]
• Bone-ash (mostly calcium phosphate) is used in the production of fine china.^[23]
• Phosphoric acid made from elemental phosphorus is used in food applications such as
  sodium tripolyphosphate.^[114] Phosphates are used to improve the characteristics of processed meat and cheese, and in toothpaste.^[114]
• smoke-screening as smoke pots and smoke bombs, and in tracer ammunition. It is also a part of an obsolete M34 White Phosphorus US hand grenade. This multipurpose grenade was mostly used for signaling, smoke screens, and inflammation; it could also cause severe burns and had a psychological impact on the enemy.^[125]
  Military uses of white phosphorus are constrained by international law.
• ³²P and ³³P are used as radioactive tracers in biochemical laboratories.[126]

Biological role

Inorganic phosphorus in the form of the phosphate PO³⁻
₄ is required for all known forms of life.^[127] 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.^[16] Biochemists commonly use the abbreviation "P_i" to refer to inorganic phosphate.^[128]

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.^[129]

An average adult human contains about 0.7 kg of phosphorus, about 85–90% in bones and teeth in the form of

; and excretion almost exclusively in the form of phosphate ions such as H
₂PO⁻
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:^[16]

Ca
₅(PO
₄)
₃OH + F⁻
→ Ca
₅(PO
₄)
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^[131]) 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.^[132]

Phosphorus is an essential

Nutrition

Dietary recommendations

The

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.^[134] The European Food Safety Authority reviewed the same safety question and decided that there was not sufficient information to set a UL.^[135]

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.^[136][137] A table of the old and new adult daily values is provided at Reference Daily Intake.

Food sources

The main food sources for phosphorus are the same as those containing protein, although proteins do not contain phosphorus. For example, milk, meat, and soya typically also have phosphorus. As a rule, if a diet has sufficient protein and calcium, the amount of phosphorus is probably sufficient.^[138]

Precautions

Organic compounds of phosphorus form a wide class of materials; many are required for life, but some are extremely toxic. Fluorophosphate

The white phosphorus allotrope presents a significant hazard because it ignites in the air and produces phosphoric acid residue. Chronic white phosphorus poisoning leads to necrosis of the jaw called "phossy jaw". White phosphorus is toxic, causing severe liver damage on ingestion and may cause a condition known as "Smoking Stool Syndrome".^[139]

In the past, external exposure to elemental phosphorus was treated by washing the affected area with 2% copper(II) sulfate solution to form harmless compounds that are then washed away. According to the recent US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries, "Cupric (copper(II)) sulfate has been used by U.S. personnel in the past and is still being used by some nations. However, copper sulfate is toxic and its use will be discontinued. Copper sulfate may produce kidney and cerebral toxicity as well as intravascular hemolysis."^[140]

The manual suggests instead "a bicarbonate solution to neutralise phosphoric acid, which will then allow removal of visible white phosphorus. Particles often can be located by their emission of smoke when air strikes them, or by their phosphorescence in the dark. In dark surroundings, fragments are seen as luminescent spots. Promptly

] As white phosphorus readily mixes with oils, any oily substances or ointments are not recommended until the area is thoroughly cleaned and all white phosphorus removed.

People can be exposed to phosphorus in the workplace by inhalation, ingestion, skin contact, and eye contact. The

US DEA List I status

Phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine.^[142] 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.^[143] In the United States, handlers of red or white phosphorus are subject to stringent regulatory controls.^{[143][144][145]}

See also

• Hubbert peak theory

Notes

Bibliography

References

General sources

• Beatty, Richard (2000). Phosphorus. Marshall Cavendish.
  ISBN 0-7614-0946-7
  .
• Burns, Melinda (10 February 2010). "The Story of P(ee)". Miller-McCune. Archived from the original on 7 January 2012. Retrieved 2 February 2012.
• Corbridge, D. E. C. (1995) "Phosphorus: An Outline of its Chemistry, Biochemistry, and Technology" 5th Edition Elsevier: Amsterdam.
  ISBN 0-444-89307-5
  .
• S2CID 1450932
  .
• Cordell, Dana & Stuart White 2011. Review: Peak Phosphorus: Clarifying the Key Issues of a Vigorous Debate about Long-Term Phosphorus Security. Sustainability 2011, 3(10), 2027-2049; doi:10.3390/su3102027, http://www.mdpi.com/2071-1050/3/10/2027/htm
• Diskowski, Herbert; Hofmann, Thomas. "Phosphorus".
  ISBN 978-3527306732
  .
• Emsley, John (2000). The Shocking history of Phosphorus: A biography of the Devil's Element. London: MacMillan.
  ISBN 0-333-76638-5
  .
• Emsley, John (7 January 2002). The 13th Element: The Sordid Tale of Murder, Fire, and Phosphorus. John Wiley & Sons.
  ISBN 978-0-471-44149-6
  . Retrieved 3 February 2012.
• Heckenmüller, M.; Narita, D.; Klepper, G. (2014). "Global availability of phosphorus and its implications for global food supply: An economic overview". Kiel Working Paper no. 1897. Accessed 11 May 2020.
• Phosphorus in diet: MedlinePlus Medical Encyclopedia. Nlm.nih.gov (2011-11-07). Retrieved on 2011-11-19.
• Parkes, G. D.; Mellor, J. W. (1939). Mellor's Modern Inorganic Chemistry. Longman's Green and Co.
• Ruttenberg, K. C. Phosphorus Cycle – Terrestrial Phosphorus Cycle, Transport of Phosphorus, from Continents to the Ocean, The Marine Phosphorus Cycle. Archived 2011-07-13 at archive.today.
• Sommers, Michael A. (2007-08-15). Phosphorus. Rosen Group.
  ISBN 978-1-4042-1960-1.{{cite book}}: CS1 maint: date and year (link
  )
• Schrödter, Klaus; Bettermann, Gerhard; Staffel, Thomas; Wahl, Friedrich; Klein, Thomas; Hofmann, Thomas. "Phosphoric Acid and Phosphates".
  ISBN 978-3527306732
  .
• Threlfall, Richard E. (1951). The Story of 100 years of Phosphorus Making: 1851–1951. Oldbury: Albright & Wilson Ltd.
• Toy, Arthur D. F. (1975). The Chemistry of Phosphorus. Texts in Inorganic Chemistry. Vol. 3. Pergamon.
  ISBN 978-1-4831-4741-3
  . Retrieved 2013-10-22.

Further reading

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