======================================================================
= Phosphorus =
======================================================================
Introduction
======================================================================
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 {{chem2|^{31}P}} as its only stable isotope,
phosphorus has an occurrence in Earth's crust of about 0.1%, generally
as phosphate rock. A member of the pnictogen family, phosphorus
readily forms a wide variety of organic and inorganic compounds, with
as its main oxidation states +5, +3 and −3.
The isolation of white phosphorus in 1669 by Hennig Brand marked the
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 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
======================================================================
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. 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 . The word phosphorus
itself () originates from Greek mythology, where it references the god
of the morning star, also known as the planet Venus.
Brand at first tried to keep the method secret, but later sold the
recipe for 200 thalers to from Dresden. 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. 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:
:
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. First, grinding up the bones into their
constituent tricalcium phosphate and treating it with sulfuric acid:
:
Then, dehydrating the resulting monocalcium phosphate:
:
Finally, mixing the obtained calcium metaphosphate with ground coal or
charcoal in an iron pot, and distilling phosphorus vapour out of a
retort:
:
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.
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 (), 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.
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).
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.
Isotopes
==========
There are 22 known isotopes of phosphorus, ranging from
{{chem2|^{26}P}} to {{chem2|^{47}P}}. Only {{chem2|^{31}P}} is stable
and is therefore present at 100% abundance. The half-integer nuclear
spin and high abundance of {{chem2|^{31}P}} 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:
* {{chem2|^{32}P|link=phosphorus-32}}, 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.
* {{chem2|^{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 {{chem2|^{32}P}} penetrate skin
and corneas and any {{chem2|^{32}P}} ingested, inhaled, or absorbed is
readily incorporated into bone and nucleic acids. For these reasons,
personnel working with {{chem2|^{32}P}} 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
============
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. Structurally, it is composed of
tetrahedra. The nature of bonding in a given 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 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 tetrahedra. The α-form is most stable at room temperature
and has a cubic crystal structure. When cooled down to 195.2 K 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 and boiling at 553.7 K, before facing thermal
decomposition at 1100 K where it turns into gaseous diphosphorus ().
This molecule contains a triple bond and is analogous to ; it can also
be generated as a transient intermediate in solution by thermolysis of
organophosphorus precursor reagents. At still higher temperatures,
dissociates into atomic P.
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 and 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 wherein one P-P bond is broken and one additional bond
is formed with the neighbouring tetrahedron, resulting in chains of
molecules linked by van der Waals forces. Red phosphorus may be formed
by heating white phosphorus to 250 C in the absence of air or by
exposing it to sunlight. 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. 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. 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. 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 12000 atm). 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.
Phosphoric acids
==================
The most prevalent compounds of phosphorus are derivatives of
phosphate (), 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:
: ('K'a1 = 7.25×10−3)
: ('K'a2 = 6.31×10−8)
: ('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 ().
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
and . For example, the industrially important pentasodium triphosphate
(also known as sodium tripolyphosphate, STPP) is produced industrially
by the megatonne by this condensation reaction:
:
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 hypophosphorous acid 1 acid, salts
+3 phosphorous acid (phosphonic acid) 2 acid, salts
+3 metaphosphorous acid 1 salts
+4 hypophosphoric acid 4 acid, salts
+5 {{chem2|(HPO3)_{'n'}}||metaphosphoric acids||'n'||salts ('n' =
3,4,6)
|-
| +5||{{chem2|H(HPO3)_{'n'}OH}}||polyphosphoric acids||'n'+2||acids,
salts ('n' = 1-6)
|-
| +5||||tripolyphosphoric acid||3||salts
|-
| +5||||pyrophosphoric acid||4||acid, salts
|-
| +5||||(ortho)phosphoric acid||3||acid, salts
|}
Oxides and sulfides
=====================
Phosphorus pentoxide () 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 (, also
called tetraphosphorus hexoxide) is the anhydride of , the minor
tautomer of phosphorous acid. The structure of is like that of
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 () has a structure analogous to ,
and is used in the manufacture of additives and pesticides. The
three-fold symmetric Phosphorus sesquisulfide () 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 , the yellowish liquids and , and the solid . 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. is
toxic because it binds to haemoglobin.
Most phosphorus pentahalides are common compounds. is a colourless
gas and the molecules have a trigonal bipyramidal geometry. With
fluoride, it forms , an anion that is isoelectronic with . is a
colourless solid which has an ionic formulation of , but adopts a
trigonal bipyramidal geometry when molten or in the vapour phase. Both
the pentafluoride and the pentachloride are Lewis acids. Meanwhile,
is an unstable solid formulated as . is not known.
The most important phosphorus oxyhalide is phosphorus oxychloride (),
which is approximately tetrahedral. It is prepared from and used in
the manufacture of plasticizers. Phosphorus can also form thiohalides
such as , and in rare cases selenohalides.
Nitrides
==========
The PN molecule phosphorus mononitride is considered unstable, but is
a product of crystalline triphosphorus pentanitride decomposition at
1100 K. Similarly, is considered unstable, and phosphorus nitride
halogens like , , , and oligomerise into cyclic polyphosphazenes. For
example, compounds of the formula {{chem2|(PNCl2)_{'n'}|}} exist
mainly as rings such as the trimer hexachlorophosphazene. The
phosphazenes arise by treatment of phosphorus pentachloride with
ammonium chloride:
:{{chem2|PCl5 + NH4Cl -> 1/'n' (NPCl2)_{'n'} + 4 HCl}}
When the chloride groups are replaced by alkoxide (), 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 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 . These compounds react with
water to form phosphine. Some phosphide minerals are also known, like
and , but they are very rare on Earth, most instances occurring in
iron-nickel meteorites.
Phosphine () and its organic derivatives are structural analogues of
ammonia (), 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 (). 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 {{chem2|P_{'n'}H_{'n'+2}|}}. The
highly flammable gas diphosphine () is an analogue of hydrazine.
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
serves as a source of in routes to organophosphorus(III) compounds.
For example, it is the precursor to triphenylphosphine:
:
Treatment of phosphorus trihalides with alcohols and phenols gives
phosphites, e.g. triphenylphosphite:
:
Similar reactions occur for phosphorus oxychloride, affording
triphenylphosphate:
:
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. ) 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.
Cells
=======
Inorganic phosphorus in the form of the phosphate 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:
:
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. 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% . 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 (),
and calcium sulfate dihydrate () 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.
Mining
========
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 1000000 ST 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 , and coke to produce . The product, being volatile,
is readily isolated:
:
:
Side products from the thermal process include ferrophosphorus, a
crude form of , 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
==========
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.
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
============================
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.
Matches
=========
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 ^{30}Si}} into {{chem2|^{31}P}} by neutron capture and beta
decay as follows:
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.
External contact
==================
{{Chembox
| container_only = yes
| Section8 =
}}
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:
Instead, the manual suggests:
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
======================================================================
{{reflist|refs=
License
=========
All content on Gopherpedia comes from Wikipedia, and is licensed under CC-BY-SA
License URL:
http://creativecommons.org/licenses/by-sa/3.0/
Original Article:
http://en.wikipedia.org/wiki/Phosphorus
