Phosphorus is a chemical element with symbol P and atomic number 15.
As an element, phosphorus exists in two major forms, white phosphorus
and red phosphorus, but because it is highly reactive, phosphorus is
never found as a free element on Earth. With a concentration of
0.099%, phosphorus is the most abundant pnictogen in the Earth's
crust. Other than a few exceptions, minerals containing phosphorus are
in the maximally oxidized state as inorganic phosphate rocks.
The first form of elemental phosphorus that was produced (white
phosphorus, in 1669) emits a faint glow when exposed to oxygen –
hence the name, taken from Greek mythology, Φωσφόρος meaning
"light-bearer" (Latin Lucifer), referring to the "Morning Star", the
Venus (or Mercury). The term "phosphorescence", meaning glow
after illumination, originally derives from this property of
phosphorus, although this word has since been used for a different
physical process that produces a glow. The glow of phosphorus itself
originates from oxidation of the white (but not red) phosphorus — a
process now termed chemiluminescence. Together with nitrogen, arsenic,
antimony, and bismuth, phosphorus is classified as a pnictogen.
Phosphorus is essential for life. Phosphates (compounds containing the
phosphate ion, PO43−) are a component of DNA, RNA, ATP, and
phospholipids. Elemental phosphorus was first isolated from human
urine, and bone ash was an important early phosphate source. Phosphate
mines contain fossils because phosphate is present in the fossilized
deposits of animal remains and excreta. Low phosphate levels are an
important limit to growth in some aquatic systems. The vast majority
of phosphorus compounds produced are consumed as fertilisers.
Phosphate is needed to replace the phosphorus that plants remove from
the soil, and its annual demand is rising nearly twice as fast as the
growth of the human population. Other applications include
organophosphorus compounds in detergents, pesticides, and nerve
2.2 Crust and organic sources
3.3 Phosphorus(I) and phosphorus(II)
3.4 Phosphides and phosphines
3.8 Organophosphorus compounds
Bone ash and guano
5.1 Peak phosphorus
5.2 Elemental phosphorus
6.3 Metallurgical aspects
6.5 Water softening
7 Biological role
Bone and teeth enamel
7.3 Dietary recommendations
7.4 Food sources
8.1 US DEA List I status
Main article: Allotropes of phosphorus
White phosphorus exposed to air glows in the dark
Crystal structure of red phosphorus
Crystal structure of black phosphorus
Phosphorus has several allotropes that exhibit strikingly different
properties. The two most common allotropes are white phosphorus and
From the perspective of applications and chemical literature, the most
important form of elemental phosphorus is white phosphorus, often
abbreviated as WP. It is a soft and waxy solid which consists of
4 molecules, in which each atom is bound to the other three atoms by a
single bond. This P
4 tetrahedron is also present in liquid and gaseous phosphorus up to
the temperature of 800 °C (1,470 °F) when it starts
decomposing to P
White phosphorus exists in two crystalline forms: α
(alpha) and β (beta). At room temperature, the α-form is stable,
which is more common and it 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.
White phosphorus is the least stable, the most reactive, the most
volatile, the least dense, and the most toxic of the allotropes. White
phosphorus gradually changes to red phosphorus. This transformation is
accelerated by light and heat, and samples of white phosphorus almost
always contain some red phosphorus and accordingly appear yellow. For
this reason, white phosphorus that is aged or otherwise impure (e.g.,
weapons-grade, not lab-grade WP) is also called yellow phosphorus.
When exposed to oxygen, white phosphorus glows in the dark with a very
faint tinge of green and blue. It is highly flammable and pyrophoric
(self-igniting) upon contact with air. Owing to its pyrophoricity,
white phosphorus is used as an additive in napalm. The odour of
combustion of this form has a characteristic garlic smell, and samples
are commonly coated with white "phosphorus pentoxide", which consists
10 tetrahedra with oxygen inserted between the phosphorus atoms and at
White phosphorus is insoluble in water but soluble in
Thermolysis of P4 at 1100 kelvin 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. At still higher temperatures, P2 dissociates
into atomic P.
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 a
Red phosphorus may be formed by heating white
phosphorus to 250 °C (482 °F) or by exposing white
phosphorus to sunlight.
Phosphorus after this treatment is
amorphous. Upon further heating, this material crystallises. In this
sense, red phosphorus is not an allotrope, but rather an intermediate
phase between the white and violet phosphorus, and most of its
properties have a range of values. For example, freshly prepared,
bright red phosphorus is highly reactive and ignites at about
300 °C (572 °F), though it is more stable than white
phosphorus, which ignites at about 30 °C (86 °F).
After prolonged heating or storage, the color darkens (see infobox
images); the resulting product is more stable and does not
spontaneously ignite in air.
Violet phosphorus is a form of phosphorus that can be produced by
day-long annealing of red phosphorus above 550 °C. In 1865,
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" (or violet or α-metallic
Black phosphorus is the least reactive allotrope and the
thermodynamically stable form below 550 °C (1,022 °F). It
is also known as β-metallic phosphorus and has a structure somewhat
resembling that of graphite. It is obtained by heating white
phosphorus under high pressures (about 12,000 standard atmospheres or
1.2 gigapascals). It can also be produced at ambient conditions using
metal salts, e.g. mercury, as catalysts. In appearance,
properties, and structure, it resembles graphite, being black and
flaky, a conductor of electricity, and has puckered sheets of linked
Another form, scarlet phosphorus, is obtained by allowing a solution
of white phosphorus in carbon disulfide to evaporate in sunlight.
Properties of some allotropes of phosphorus
Band gap (eV)
It was known from early times that the green glow emanating from white
phosphorus would persist for a time in a stoppered jar, but then
Robert Boyle in the 1680s ascribed it to "debilitation" of the
air; in fact, it is oxygen being consumed. By the 18th century, it was
known that in pure oxygen, phosphorus does not glow at all; there
is only a range of partial pressures at which it does. Heat can be
applied to drive the reaction at higher pressures.
In 1974, the glow was explained by R. J. van Zee and A. U.
Khan. A reaction with oxygen takes place at the surface of the
solid (or liquid) phosphorus, forming the short-lived molecules HPO
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 that time, 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
Isotopes of phosphorus
Twenty-three isotopes of phosphorus are known, including all
possibilities from 24P up to 46P. Only 31P is stable and is therefore
present at 100% abundance. The half-integer nuclear spin and high
abundance of 31P make phosphorus-31 NMR spectroscopy a very useful
analytical tool in studies of phosphorus-containing samples.
Two radioactive isotopes of phosphorus have half-lives suitable for
biological scientific experiments. 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
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
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, Occupational Safety and
Health Administration in the United States, and similar institutions
in other developed countries require personnel working with 32P 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.
See also: Category:
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.
Crust and organic sources
At 0.099%, phosphorus is the most abundant pnictogen in the Earth's
crust but it is not found free in nature; it is widely distributed in
many minerals, mainly phosphates. Inorganic phosphate rock, which
is partially made of apatite (a group of minerals being, generally,
pentacalcium triorthophosphate fluoride (hydroxide)), is today the
chief commercial source of this element. According to the US
Geological Survey (USGS), about 50 percent of the global phosphorus
reserves are in the Arab nations. Large deposits of apatite are
located in China, Russia, Morocco, Florida, Idaho, Tennessee,
Utah, and elsewhere.
Albright and Wilson in the UK and their
Niagara Falls plant, for instance, were using phosphate rock in the
1890s and 1900s from Tennessee, Florida, and the Îles du Connétable
(guano island sources of phosphate); by 1950, they were using
phosphate rock mainly from
Tennessee and North Africa.
Organic sources, namely urine, bone ash and (in the latter 19th
century) guano, were historically of importance but had only limited
commercial success. As urine contains phosphorus, it has
fertilising qualities which are still harnessed today in some
countries, including Sweden, using methods for reuse of excreta. To
this end, urine can be used as a fertiliser in its pure form or part
of being mixed with water in the form of sewage or sewage sludge.
See also: Category:
The tetrahedral structure of P4O10 and P4S10.
The most prevalent compounds of phosphorus are derivatives of
phosphate (PO43−), a tetrahedral anion.
Phosphate is the
conjugate base of phosphoric acid, which is produced on a massive
scale for use in fertilisers. Being triprotic, phosphoric acid
converts stepwise to three conjugate bases:
H3PO4 + H2O ⇌ H3O+ + H2PO4− Ka1=
H2PO4− + H2O ⇌ H3O+ + HPO42− Ka2=
HPO42− + H2O ⇌ H3O+ + PO43−
Phosphate exhibits the tendency to form chains and rings with P-O-P
bonds. Many polyphosphates are known, including ATP. Polyphosphates
arise by dehydration of hydrogen phosphates such as HPO42− and
H2PO4−. For example, the industrially important trisodium
triphosphate (also known as sodium tripolyphosphate, STPP) is produced
industrially on by the megatonne by this condensation reaction:
2 Na2[(HO)PO3] + Na[(HO)2PO2] → Na5[O3P-O-P(O)2-O-PO3] + 2 H2O
Phosphorus pentoxide (P4O10) is the acid anhydride of phosphoric acid,
but several intermediates between the two are known. This waxy white
solid reacts vigorously with water.
With metal cations, phosphate forms a variety of salts. These solids
are polymeric, featuring P-O-M linkages. When the metal cation has a
charge of 2+ or 3+, the salts are generally insoluble, hence they
exist as common minerals. Many phosphate salts are derived from
hydrogen phosphate (HPO42−).
PCl5 and PF5 are common compounds. PF5 is a colourless gas and the
molecules have trigonal bypramidal geometry. PCl5 is a colourless
solid which has an ionic formulation of PCl4+ PCl6−, but adopts the
trigonal bypramidal geometry when molten or in the vapour phase.
PBr5 is an unstable solid formulated as PBr4+Br−and PI5 is not
known. The pentachloride and pentafluoride are Lewis acids. With
fluoride, PF5 forms PF6−, an anion that is isoelectronic with SF6.
The most important oxyhalide is phosphorus oxychloride, (POCl3), which
is approximately tetrahedral.
Before extensive computer calculations were feasible, it was thought
that bonding in phosphorus(V) compounds involved d orbitals. Computer
modeling of molecular orbital theory indicates that this bonding
involves only s- and p-orbitals.
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
P4 + 6 Cl2 → 4 PCl3
The trifluoride is produced from the trichloride by halide exchange.
PF3 is toxic because it binds to haemoglobin.
Phosphorus(III) oxide, P4O6 (also called tetraphosphorus hexoxide) is
the anhydride of P(OH)3, the minor tautomer of phosphorous acid. The
structure of P4O6 is like that of P4O10 without the terminal oxide
Phosphorus(I) and phosphorus(II)
A stable diphosphene, a derivative of phosphorus(I).
These compounds generally feature P-P bonds. Examples include
catenated derivatives of phosphine and organophosphines. Compounds
containing P=P double bonds have also been observed, although they are
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.
Schreibersite is a naturally occurring metal rich phosphide found in
meteorites. The structures of the metal rich and phosphorus rich
phosphides can be structurally complex.
Phosphine (PH3) and its organic derivatives (PR3) are structural
analogues with ammonia (NH3) but the bond angles at phosphorus are
closer to 90° for phosphine and its organic derivatives. It is an
ill-smelling, toxic compound.
Phosphorus has an oxidation number of -3
Phosphine is produced by hydrolysis of calcium
phosphide, Ca3P2. Unlike ammonia, phosphine is oxidised by air.
Phosphine is also far less basic than ammonia. Other phophines are
known which contain chains of up to nine phosphorus atoms and have the
formula PnHn+2. The highly flammable gas diphosphine (P2H4) is an
analogue of hydrazine.
Phosphorous 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
important, and three of them, hypophosphorous acid, phosphorous acid,
and phosphoric acid, are particularly important.
acids, salts (n=1-6)
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.
Main article: phosphorus sulfide
Phosphorus forms a wide range of sulfides, where the phosphorus can be
in P(V), P(III) or other oxidation states. The most famous is the
three-fold symmetric P4S3 which is used in strike-anywhere matches.
P4S10 and P4O10 have analogous structures. Mixed oxyhalides and
oxyhydrides of phosphorus(III) are almost unknown.
Main article: organophosphorus compounds
Compounds with P-C and P-O-C bonds are often classified as
organophosphorus compounds. They are widely used commercially. The
PCl3 serves as a source of P3+ in routes to organophosphorus(III)
compounds. For example, it is the precursor to triphenylphosphine:
PCl3 + 6 Na + 3 C6H5Cl → P(C6H5)3 + 6 NaCl
Treatment of phosphorus trihalides with alcohols and phenols gives
phosphites, e.g. triphenylphosphite:
PCl3 + 3 C6H5OH → P(OC6H5)3 + 3 HCl
Similar reactions occur for phosphorus oxychloride, affording
OPCl3 + 3 C6H5OH → OP(OC6H5)3 + 3 HCl
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
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
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
The discovery of phosphorus, the first element to be discovered that
was not known since ancient times, is credited to the German
Hennig Brand in 1669, although other chemists might have
discovered phosphorus around the same time. Brand experimented
with urine, which contains considerable quantities of dissolved
phosphates from normal metabolism. 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
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. We now know that Brand produced ammonium sodium
hydrogen phosphate, (NH
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. Later
scientists discovered that fresh urine yielded the same amount of
Brand at first tried to keep the method secret, but later sold the
recipe for 200 thalers to D. Krafft from Dresden, who could now
make it as well, and toured much of Europe with it, including England,
where he met with Robert Boyle. The secret that it was made from urine
leaked out and first Johann Kunckel (1630–1703) in
Sweden (1678) and
later Boyle in London (1680) also managed to make phosphorus, possibly
with the aid of his assistant, Ambrose Godfrey-Hanckwitz, who later
made a business of the manufacture of phosphorus.
Boyle states that Krafft gave him no information as to the preparation
of phosphorus other than that it was derived from "somewhat that
belonged to the body of man". This gave Boyle a valuable clue, so that
he, too, managed to make phosphorus, and published the method of its
manufacture. Later he improved Brand's process by using sand in
the reaction (still using urine as base material),
3 + 2 SiO
2 + 10 C → 2 Na
3 + 10 CO + P
Robert Boyle was the first to use phosphorus to ignite sulfur-tipped
wooden splints, forerunners of our modern matches, in 1680.
Phosphorus was the 13th element to be discovered. For this reason, and
due to its use in explosives, poisons and nerve agents, it is
sometimes referred to as "the Devil's element".
Bone ash and guano
Guano mining in the Central Chincha Islands, ca. 1860.
Johan Gottlieb Gahn
Johan Gottlieb Gahn and
Carl Wilhelm Scheele
Carl Wilhelm Scheele showed that
calcium phosphate (Ca
2) is found in bones, and they obtained elemental phosphorus from bone
Antoine Lavoisier recognised phosphorus as an element in
Bone ash was the major source of phosphorus until the 1840s.
The method started by roasting bones, then employed the use of clay
retorts encased in a very hot brick furnace to distill out the highly
toxic elemental phosphorus product. Alternately, precipitated
phosphates could be made from ground-up bones that had been de-greased
and treated with strong acids.
White phosphorus could then be made by
heating the precipitated phosphates, mixed with ground coal or
charcoal in an iron pot, and distilling off phosphorus vapour in a
Carbon monoxide and other flammable gases produced during
the reduction process were burnt off in a flare stack.
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.
Phosphate rock, which usually contains calcium phosphate, was first
used in 1850 to make phosphorus, and following the introduction of the
electric arc furnace by
James Burgess Readman in 1888 (patented
1889), elemental phosphorus production switched from the bone-ash
heating, to electric arc production from phosphate rock. After the
depletion of world guano sources about the same time, mineral
phosphates became the major source of phosphate fertiliser production.
Phosphate rock production greatly increased after World War II, and
remains the primary global source of phosphorus and phosphorus
chemicals today. See the article on peak phosphorus for more
information on the history and present state of phosphate mining.
Phosphate rock remains a feedstock in the fertiliser industry, where
it is treated with sulfuric acid to produce various "superphosphate"
White phosphorus was first made commercially in the 19th century for
the match industry. This used bone ash for a phosphate source, as
described above. The bone-ash process became obsolete when the
submerged-arc furnace for phosphorus production was introduced to
reduce phosphate rock. The electric furnace method allowed
production to increase to the point where phosphorus could be used in
weapons of war. In World War I, it was used in incendiaries,
smoke screens and tracer bullets. A special incendiary bullet was
developed to shoot at hydrogen-filled Zeppelins over Britain (hydrogen
being highly flammable). During World War II, Molotov cocktails
made of phosphorus dissolved in petrol were distributed in Britain to
specially selected civilians within the British resistance operation,
for defence; and phosphorus incendiary bombs were used in war on a
large scale. Burning phosphorus is difficult to extinguish and if it
splashes onto human skin it has horrific effects.
Early matches used white phosphorus in their composition, which was
dangerous due to its toxicity. Murders, suicides and accidental
poisonings resulted from its use. (An apocryphal tale tells of a woman
attempting to murder her husband with white phosphorus in his food,
which was detected by the stew's giving off luminous steam). In
addition, exposure to the vapours gave match workers a severe necrosis
of the bones of the jaw, the infamous "phossy jaw". When a safe
process for manufacturing red phosphorus was discovered, with its far
lower flammability and toxicity, laws were enacted, under the Berne
Convention (1906), requiring its adoption as a safer alternative for
match manufacture. The toxicity of white phosphorus led to
discontinuation of its use in matches. The Allies used phosphorus
incendiary bombs in
World War II
World War II to destroy Hamburg, the place where
the "miraculous bearer of light" was first discovered.
Mining of phosphate rock in Nauru
Most production of phosphorus-bearing material is for agriculture
fertilisers. For this purpose, phosphate minerals are converted to
phosphoric acid. It follows two distinct chemical routes, the main one
being treatment of phosphate minerals with sulfuric acid. The other
process utilises white phosphorus, which may be produced by reaction
and distillation from very low grade phosphate sources. The white
phosphorus is then oxidised to phosphoric acid and subsequently
neutralised with base to give phosphate salts. Phosphoric acid
produced from white phosphorus is relatively pure and is the main
route for the production of phosphates for all purposes, including
In the early 1990s, Albright and Wilson's purified wet phosphoric acid
business was being adversely affected by phosphate rock sales by China
and the entry of their long-standing Moroccan phosphate suppliers into
the purified wet phosphoric acid business.
Main article: Peak phosphorus
In 2017, the USGS estimated 68 billion tons of world reserves, where
reserve figures refer to the amount assumed recoverable at current
market prices; 0.261 billion tons were mined in 2016. Critical to
contemporary agriculture, its annual demand is rising nearly twice as
fast as the growth of the human population.
The production of phosphorus may have peaked already (as per 2011),
leading to the possibility of global shortages by 2040. In 2007,
at the rate of consumption, the supply of phosphorus was estimated to
run out in 345 years. However, some scientists now believe that a
"peak phosphorus" will occur in 30 years and that "At current rates,
reserves will be depleted in the next 50 to 100 years." Cofounder
of Boston-based investment firm and environmental foundation Jeremy
Grantham wrote in Nature in November 2012 that consumption of the
element "must be drastically reduced in the next 20-40 years or we
will begin to starve." According to N.N. Greenwood and A.
Earnshaw, authors of the textbook, Chemistry of the Elements, however,
phosphorus comprises about 0.1% by mass of the average rock, and
consequently the Earth's supply is vast, although dilute.
Presently, about 1,000,000 short tons (910,000 t) of elemental
phosphorus is produced annually.
Calcium phosphate (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 (refined coal) to produce vaporised P
4. The product is subsequently condensed into a white powder under
water to prevent oxidation by air. Even under water, white phosphorus
is slowly converted to the more stable red phosphorus allotrope. The
chemical equation for this process when starting with fluoroapatite, a
common phosphate mineral, is:
4 Ca5(PO4)3F + 18 SiO2 + 30 C → 3 P4 + 30 CO + 18 CaSiO3 + 2 CaF2
Side products from this 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; train
derailments at Brownston, Nebraska and
Miamisburg, Ohio led to large
fires. The worst incident in recent times was an environmental
contamination in 1968 when the sea was polluted from spillage and/or
inadequately treated sewage from a white phosphorus plant at Placentia
Another process by which elemental phosphorus is extracted includes
applying at high temperatures (1500 °C):
2 Ca3(PO4)2 + 6 SiO2 + 10 C → 6 CaSiO3 + 10 CO + P4
Historically, before the development of mineral-based extractions,
white phosphorus was isolated on an industrial scale from bone
ash. 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
Ca(H2PO4)2 → Ca(PO3)2 + 2 H2O
When ignited to a white heat 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
Main article: Fertiliser
Phosphorus is an essential plant nutrient (often the limiting
nutrient), and the bulk of all phosphorus production is in
concentrated phosphoric acids for agriculture fertilisers, containing
as much as 70% to 75% P2O5. Its annual demand is rising nearly twice
as fast as the growth of the human population. That led to large
increase in phosphate (PO43−) production in the second half of the
20th century. Artificial phosphate fertilisation is necessary
because phosphorus is essential to all life organisms, natural
phosphorus-bearing compounds are mostly insoluble and inaccessible to
plants, and the natural cycle of phosphorus is very slow. 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
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.
Widely used compounds
Baking powder and fertilisers
Animal food additive, toothpowder
Manufacture of phosphate fertilisers
Manufacture of POCl3 and pesticides
Manufacture of plasticiser
Manufacturing of additives and pesticides
White phosphorus is widely used to make organophosphorus compounds
through intermediate phosphorus chlorides and two phosphorus sulfides,
phosphorus pentasulfide and phosphorus sesquisulfide.
Organophosphorus compounds have many applications, including in
plasticisers, flame retardants, pesticides, extraction agents, nerve
agents and water treatment.
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.
Match striking surface made of a mixture of red phosphorus, glue and
ground glass. The glass powder is used to increase the friction.
Main article: Match
The first striking match with a phosphorus head was invented by
Charles Sauria in 1830. These matches (and subsequent modifications)
were made with heads of white phosphorus, an oxygen-releasing compound
(potassium chlorate, lead dioxide, or sometimes nitrate), and a
binder. They were poisonous to the workers in manufacture,
sensitive to storage conditions, toxic if ingested, and hazardous when
accidentally ignited on a rough surface. Production in several
countries was banned between 1872 and 1925. The international
Berne Convention, ratified in 1906, prohibited the use of white
phosphorus in matches.
In consequence, the 'strike-anywhere' matches were gradually replaced
by 'safety matches', wherein the white phosphorus was replaced by
phosphorus sesquisulfide (P4S3), sulfur, or antimony sulfide. Such
matches are difficult to ignite on any surface other than a special
strip. The strip contains red phosphorus that heats up upon striking,
reacts with the oxygen-releasing compound in the head, and ignites the
flammable material of the head.
Sodium tripolyphosphate made from phosphoric acid is used in laundry
detergents in some countries, but banned for this use in others.
This compound softens the water to enhance the performance of the
detergents and to prevent pipe/boiler tube corrosion.
Phosphates are used to make special glasses for sodium lamps.
Bone-ash, calcium phosphate, is used in the production of fine
Phosphoric acid made from elemental phosphorus is used in food
applications such as soft drinks, and as a starting point for food
grade phosphates. These include mono-calcium phosphate for baking
powder and sodium tripolyphosphate. Phosphates are used to improve
the characteristics of processed meat and cheese, and in
White phosphorus, called "WP" (slang term "Willie Peter") is used in
military applications as incendiary bombs, for 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.
Military uses of white phosphorus are
constrained by international law.
In trace amounts, phosphorus is used as a dopant for n-type
32P and 33P are used as radioactive tracers in biochemical
Phosphate is a strong complexing agent for the hexavalent uranyl
(UO22+) species; for this reason, apatite and other natural phosphates
can be very rich in uranium.
Tributylphosphate is an organophosphate soluble in kerosene used to
extract uranium in the
Purex process for reprocessing spent nuclear
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.
Phospholipids are the main structural components of all cellular
Calcium phosphate salts assist in stiffening bones.
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.
An average adult human contains about 0.7 kg of phosphorus, about
85–90% in bones and teeth in the form of apatite, and the remainder
in soft tissues and extracellular fluids (~1%). The phosphorus content
increases from about 0.5 weight% in infancy to 0.65–1.1 weight% in
adults. Average phosphorus concentration in the blood is about 0.4
g/L, about 70% of that is organic and 30% inorganic phosphates. An
adult with healthy diet consumes and excretes about 1–3 grams
of phosphorus per day, with consumption in the form of inorganic
phosphate and phosphorus-containing biomolecules such as nucleic acids
and phospholipids; and excretion almost exclusively in the form of
phosphate ions such as H
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 called
3OH + F− → Ca
3F + OH−
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) or pass 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 is an essential macromineral for plants, which is studied
extensively in edaphology to understand plant uptake from soil
Phosphorus is a limiting factor in many ecosystems; that is,
the scarcity of phosphorus limits the rate of organism growth. An
excess of phosphorus can also be problematic, especially in aquatic
systems where eutrophication sometimes leads to algal blooms.
Institute of Medicine
Institute of Medicine (IOM) updated Estimated Average
Requirements (EARs) and Recommended Dietary Allowances (RDAs) for
phosphorus in 1997. If there is not sufficient information to
establish EARs and RDAs, an estimate designated
Adequate Intake (AI)
is used instead. The current EAR for phosphorus for people ages 19 and
up is 580 mg/day. The RDA is 700 mg/day. RDAs are higher than EARs so
as to identify amounts that will cover people with higher than average
requirements. RDA for pregnancy and lactation are also 700 mg/day. For
children ages 1–18 years the RDA increases with age from 460 to 1250
mg/day. As for safety, the IOM sets Tolerable upper intake levels
(ULs) for vitamins and minerals when evidence is sufficient. In the
case of phosphorus the UL is 4000 mg/day. Collectively the EARs, RDAs,
AIs and ULs are referred to as Dietary Reference Intakes (DRIs).
European Food Safety Authority
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 defined the same as in 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 years the AI is 440 mg/day, for ages 11-17 640
mg/day. These AIs are lower than the U.S RDAs. In both systems,
teenagers need more than adults. The European Food Safety
Authority reviewed the same safety question and decided that there was
not sufficient information to set a UL.
For U.S. food and dietary supplement labeling purposes the amount in a
serving is expressed as a percent of Daily Value (%DV). For phosphorus
labeling purposes 100% of the Daily Value was 1000 mg, but as of May
27, 2016 it was revised to 1250 mg to bring it into agreement with the
RDA. A table of the old and new adult Daily Values is provided at
Reference Daily Intake. The original deadline to be in compliance was
July 28, 2018, but on September 29, 2017 the FDA released a proposed
rule that extended the deadline to January 1, 2020 for large companies
and January 1, 2021 for small companies.
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
Organic compounds of phosphorus form a wide class of materials; many
are required for life, but some are extremely toxic. Fluorophosphate
esters are among the most potent neurotoxins known. A wide range of
organophosphorus compounds are used for their toxicity as pesticides
(herbicides, insecticides, fungicides, etc.) and weaponised as nerve
agents against enemy humans. Most inorganic phosphates are relatively
nontoxic and essential nutrients.
The white phosphorus allotrope presents a significant hazard because
it ignites in 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".
In the past, external exposure to elemental phosphorus was treated by
washing the affected area with 2% copper 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
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
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."[note
1] 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 Occupational Safety and
Health Administration (OSHA) has set the phosphorus exposure limit
(Permissible exposure limit) in the workplace at 0.1 mg/m3 over
an 8-hour workday. The National Institute for Occupational Safety and
Health (NIOSH) has set a
Recommended exposure limit (REL) of
0.1 mg/m3 over an 8-hour workday. At levels of 5 mg/m3,
phosphorus is immediately dangerous to life and health.
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. For this reason, red and white phosphorus were
designated by the United States
Drug Enforcement Administration
Drug Enforcement Administration as
List I precursor chemicals under 21 CFR 1310.02 effective on November
17, 2001. In the United States, handlers of red or white
phosphorus are subject to stringent regulatory controls.
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and if there is any WP in the wound, covered by tissue or fluids such
as blood serum, it will not chemoluminesce until it is moved to a
position where the air can get at it and activate the chemoluminescent
glow, which requires a very dark room and dark-adapted eyes to see
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Periodic table (Large cells)
Alkaline earth metal
BNF: cb119329562 (data)