Phosphorus
Phosphorus is a nonmetallic chemical element
with symbol P and atomic number 15. A multivalent
pnictogen, phosphorus as a mineral is almost
always present in its maximally oxidised state,
as inorganic phosphate rocks. Elemental phosphorus
exists in two major forms—white phosphorus
and red phosphorus—but due to its high reactivity,
phosphorus is never found as a free element
on Earth.
The first form of elemental phosphorus to
be produced (white phosphorus, in 1669) emits
a faint glow upon exposure to oxygen – hence
its name given from Greek mythology, Φωσφόρος
meaning "light-bearer" (Latin Lucifer), referring
to the "Morning Star", the planet Venus. 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.
The vast majority of phosphorus compounds
are consumed as fertilisers. Other applications
include the role of organophosphorus compounds
in detergents, pesticides and nerve agents,
and matches.
Phosphorus is essential for life. As phosphate,
it is a component of DNA, RNA, ATP, and also
the phospholipids that form all cell membranes.
Demonstrating the link between phosphorus
and life, elemental phosphorus was historically
first isolated from human urine, and bone
ash was an important early phosphate source.
Phosphate minerals are fossils. Low phosphate
levels are an important limit to growth in
some aquatic systems. The chief commercial
use of phosphorus compounds for production
of fertilisers is due to the need 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.
Characteristics
Physical
Phosphorus exists as several forms (allotropes)
that exhibit strikingly different properties.
The two most common allotropes are white phosphorus
and red phosphorus. Another form, scarlet
phosphorus, is obtained by allowing a solution
of white phosphorus in carbon disulfide to
evaporate in sunlight. Black phosphorus is
obtained by heating white phosphorus under
high pressures (about 12,000 standard atmospheres
or 1.2 gigapascals). In appearance, properties,
and structure, it resembles graphite, being
black and flaky, a conductor of electricity,
and has puckered sheets of linked atoms. Another
allotrope is diphosphorus; it contains a phosphorus
dimer as a structural unit and is highly reactive.
White phosphorus and related molecular forms
The most important form of elemental phosphorus
from the perspective of applications and the
chemical literature is white phosphorus, often
abbreviated as WP. It consists of tetrahedral
P
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 when it starts decomposing to P
2 molecules. Solid white exists in two forms.
At low-temperatures, the β form is stable.
At high-temperatures the α form is predominant.
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. White phosphorus
glows in the dark (when exposed to oxygen)
with a very faint tinge of green and blue,
is highly flammable and pyrophoric (self-igniting)
upon contact with air and is toxic (causing
severe liver damage on ingestion). 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
"(di)phosphorus pentoxide", which consists
of P
4O
10 tetrahedra with oxygen inserted between
the phosphorus atoms and at their vertices.
White phosphorus is insoluble in water but
soluble in carbon disulfide.
Thermolysis (cracking) 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.
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).
Red phosphorus
Red phosphorus is polymeric in structure.
It can be viewed as a derivative of P4 wherein
one P-P bond is broken, and one additional
bond is formed with the neighbouring tetrahedron
resulting in a chain-like structure. 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, though
it is still more stable than white phosphorus,
which ignites at about 30 °C. 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
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
recrystallized from molten lead, a red/purple
form is obtained. Therefore this form is sometimes
known as "Hittorf's phosphorus" (or violet
or α-metallic phosphorus).
Black phosphorus
Black phosphorus is the least reactive allotrope
and the thermodynamically stable form below
550 °C. It is also known as β-metallic
phosphorus and has a structure somewhat resembling
that of graphite. High pressures are usually
required to produce black phosphorus, but
it can also be produced at ambient conditions
using metal salts as catalysts.
Isotopes
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 produce radiolabeled
DNA and RNA probes, e.g. for use in Northern
blots or Southern blots. Because the high
energy beta particles produced penetrate skin
and corneas, and because any 32P ingested,
inhaled, or absorbed is readily incorporated
into bone and nucleic acids, Occupational
Safety and Health Administration in the United
States, and similar institutions in other
developed countries require that a lab coat,
disposable gloves and safety glasses or goggles
be worn when working with 32P, and that working
directly over an open container be avoided
in order to protect the eyes. Monitoring personal,
clothing, and surface contamination is also
required. In addition, due to the high energy
of the beta particles, shielding this radiation
with the normally used dense materials (e.g.
lead), gives rise to secondary emission of
X-rays via Bremsstrahlung (braking radiation).
Therefore shielding must be accomplished with
low density materials, e.g. Plexiglas (Lucite),
other plastics, water, or (when transparency
is not required), even wood.
33P, a beta-emitter (0.25 MeV) with a half-life
of 25.4 days. It is used in life-science laboratories
in applications in which lower energy beta
emissions are advantageous such as DNA sequencing.
Occurrence
Phosphorus is not found free in nature, but
it is widely distributed in many minerals,
mainly phosphates. Historically-important
but limited commercial sources were organic,
such as bone ash and (in the latter 19th century)
guano. Inorganic phosphate rock, which is
partially made of apatite (an impure tri-calcium
phosphate mineral), 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. 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.
In 2012, the USGS estimated 71 billion tons
of world reserves, where reserve figures refer
to the amount assumed recoverable at current
market prices; 0.19 billion tons were mined
in 2011. Critical to contemporary agriculture,
its annual demand is rising nearly twice as
fast as the growth of the human population.
Recent reports suggest that production of
phosphorus may have peaked, 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. This book, however, was published
in 1997.
Research into urine diversion
Urine contains most (94% according to Wolgast)
of the NPK nutrients excreted by the human
body. Conversely, concentrations of heavy
metals such as lead, mercury, and cadmium,
commonly found in solid human waste, are much
lower in urine (although not low enough to
qualify for use in organic agriculture under
current EU rules). The more general limitations
to using urine as fertilizer then depend mainly
on the potential for buildup of excess nitrogen
(due to the high ratio of that macronutrient),
and inorganic salts such as sodium chloride,
which are also part of the wastes excreted
by the renal system. The degree to which these
factors impact the effectiveness depends on
the term of use, salinity tolerance of the
plant, soil composition, addition of other
fertilizing compounds, and quantity of rainfall
or other irrigation.
Urine typically contains 70% of the nitrogen
and more than half the phosphorus and potassium
found in urban waste water flows, while making
up less than 1% of the overall volume. Thus
far, source separation, or urine diversion
and on-site treatment has been implemented
in South Africa, China, and Sweden among other
countries with the Bill and Melinda Gates
Foundation provided some of the funding implementations.
China reportedly had 685,000 operating source
separation toilets spread out among 17 provinces
in 2003.
"Urine management" is a relatively new way
to view closing the cycle of agricultural
nutrient flows and reducing sewage treatment
costs and ecological consequences such as
eutrophication resulting from the influx of
nutrient rich effluent into aquatic or marine
ecosystems. Proponents of urine as a natural
source of agricultural fertilizer claim the
risks to be negligible or acceptable. Their
views seem to be backed by research showing
there are more environmental problems when
it is treated and disposed of compared with
when it is used as a resource.
It is unclear whether source separation, urine
diversion, and on-site urine treatment can
be made cost effective; nor whether required
behavioral changes would be regarded as socially
acceptable, as the largely successful trials
performed in Sweden may not readily generalize
to other industrialized societies. In developing
countries the use of whole raw sewage (night
soil) has been common throughout history,
yet the application of pure urine to crops
is rare. Increasingly there are calls for
urine's use as a fertilizer, such as a 2010
Scientific American article "Human urine is
an effective fertilizer".
Production
The majority of phosphorus-containing compounds
are produced for use as fertilisers. For this
purpose, phosphate-containing minerals are
converted to phosphoric acid. Two distinct
routes are employed, 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 obtained
via white phosphorus is relatively pure and
is the main source of phosphates used in detergents
and other non-fertiliser applications.
Elemental phosphorus
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 (impure carbon) to produce vaporized
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:
Side products from this production 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
one in 1968 when the sea became contaminated
due to spillages and/or inadequately treated
sewage from a white phosphorus plant at Placentia
Bay, Newfoundland.
Another process by which elemental phosphorus
is extracted includes applying at high temperatures
(1500 °C):
Thermphos International is Europe's only producer
of elemental phosphorus with the annual capacity
of 80,000 t provided by a plant at Vlissingen,
Netherlands.
Historically, before technology allowed mineral-based
extractions, white phosphorus was isolated
on an industrial scale from bone ash. In this
scheme the tricalcium phosphate in bone ash
is converted to monocalcium phosphate with
sulfuric acid:
Monocalcium phosphate is then dehydrated to
the corresponding metaphosphate:
Calcium metaphosphate yields, when ignited
to a white heat with charcoal, two-thirds
of its weight of white phosphorus while one-third
of the phosphorus remains in the residue as
calcium orthophosphate:
Compounds
Oxoacids of phosphorus
Phosphorous oxoacids are extensive, often
commercially important, and sometimes structurally
complicated. They all have acidic protons
bound to oxygen atoms and some have nonacidic
protons that are bonded directly to phosphorus.
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
ones.
Phosphorus(V) compounds
Oxides
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:
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 a megatonne scale via this
condensation reaction:
Phosphorus pentoxide (P4O10) is the acid anhydride
of phosphoric acid, but several intermediates
are known between the two. 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. Both are
volatile and pale or colourless. The other
two halides, PBr5 and PI5PI5 are unstable.
The pentachloride and pentafluoride adopt
trigonal bipyramid molecular geometry and
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 tetrahedral.
Before extensive computer calculations were
feasible, it was proposed that bonding in
phosphorus(V) compounds involved d orbitals.
It is now accepted that the bonding can be
better explained by molecular orbital theory
and involves only s- and p-orbitals on phosphorus.
Nitrides
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 The chloride
groups can be replaced by alkoxide (RO-) to
give rise to a family of polymers with potentially
useful properties.
Sulfides
Phosphorus forms a wide range of sulfides,
where phosphorus can be P(V), P(III) or other
oxidation states. Most famous is the three-fold
symmetric P4S3 used in strike-anywhere matches.
P4S10 and P4O10 have analogous structures.
Phosphorus(III) compounds
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. Phosphine is produced by hydrolysis
of calcium phosphide, Ca3P2. Unlike ammonia,
phosphine is oxidised by air. Phosphine is
also far less basic than ammonia.
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 phosphorus
acid. The trichloride, a common reagent, is
produced by chlorination of white phosphorus:
The trifluoride is produced by 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 less the terminal
oxide groups.
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:
Treatment of phosphorus trihalides with alcohols
and phenols gives phosphites, e.g. triphenylphosphite:
Similar reactions occur for phosphorus oxychloride,
affording triphenylphosphate:
Phosphorus(I) and phosphorus(II) compounds
These compounds generally feature P-P bonds.
Examples include catenated derivatives of
phosphine and organophosphines. The highly
flammable gas diphosphine (P2H4) is the first
of a series of derivatives of this type. Diphosphine
is an analogue of hydrazine. Compounds containing
P=P double bonds have also been observed,
although they are rare.
Phosphides
The phosphide ion is P3-. Phosphides arise
by reaction of metals with red phosphorus.
Salts of P3- do not exist in solution and
these derivatives are refractory, reflecting
their high lattice energy. Illustrated by
the behaviour calcium phosphide, many metal
phosphides hydrolyse in water with release
of phosphine:
Schreibersite is a naturally occurring phosphide
found in meteorites. Many polyphosphides are
known such as derivatives of OsP2. These can
be structurally complex ranging from Na3P7
and derivatives of P264-. Often these species
adopt cage-like structures that resemble fragments
of violet phosphorus.
Spelling and etymology
The name Phosphorus in Ancient Greece was
the name for the planet Venus and is derived
from the Greek words (φως = light, φέρω
= carry), which roughly translates as light-bringer
or light carrier. (In Greek mythology and
tradition, Augerinus (Αυγερινός = morning
star, in use until today), Hesperus or Hesperinus
(΄Εσπερος or Εσπερινός or
Αποσπερίτης = evening star, in
use until today) and Eosphorus (Εωσφόρος
= dawnbearer, not in use for the planet after
Christianity) are close homologues, and also
associated with Phosphorus-the-planet).
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).
History and discovery
Phosphorus was the 13th element to be discovered.
For this reason, and also due to its use in
explosives, poisons and nerve agents, it is
sometimes referred to as "the Devil's element".
It was the first element to be discovered
that was not known since ancient times. The
discovery of phosphorus is credited to the
German alchemist 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 light"). His 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)NaHPO
4. While the quantities were essentially correct
(it took about 1,100 litres 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 phosphorus.
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),
Robert Boyle was the first to use phosphorus
to ignite sulfur-tipped wooden splints, forerunners
of our modern matches, in 1680.
In 1769 Johan Gottlieb Gahn and Carl Wilhelm
Scheele showed that calcium phosphate (Ca
3(PO
4)
2) is found in bones, and they obtained elemental
phosphorus from bone ash. Antoine Lavoisier
recognized phosphorus as an element in 1777.
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 be 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
retort. 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 fertilizer in the latter
half of the 19th century.
Phosphate rock, a mineral containing calcium
phosphate, was first used in 1850 to make
phosphorus, and following the introduction
of the electric arc furnace in 1890, 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 fertilizer 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 fertilizer industry,
where it is treated with sulfuric acid to
produce various "superphosphate" fertilizer
products.
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 to destroy Hamburg,
the place where the "miraculous bearer of
light" was first discovered.
Glow
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 cease. 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 pressure 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 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 that time, phosphors and phosphorescence
were used loosely to describe substances that
shine in the dark without burning. Although
the term phosphorescence was originally coined
as a term by analogy with the glow from oxidation
of elemental phosphorus, it is now reserved
for another fundamentally different process—re-emission
of light after illumination.
Applications
Fertiliser
The dominant application of phosphorus is
in fertilisers, which provides phosphate as
required for all life and is often a limiting
nutrient for crops. Phosphorus, being an essential
plant nutrient, finds its major use as a constituent
of fertilizers for agriculture and farm production
in the form of concentrated phosphoric acids,
which can consist of 70% to 75% P2O5. Global
demand for fertilisers led to large increase
in phosphate (PO43–) production in the second
half of the 20th century. Due to the essential
nature of phosphorus to living organisms,
the low solubility of natural phosphorus-containing
compounds, and the slow natural cycle of phosphorus,
the agricultural industry relies on fertilisers
that contain phosphate. A major form of these
fertilisers is superphosphate of lime, a mixture
of two salts, calcium dihydrogen phosphate
Ca(H2PO4)2 and calcium sulfate dihydrate CaSO4·2H2O,
produced by the reaction of sulfuric acid
and water with calcium phosphate.
Organophosphorus compounds
White phosphorus is widely used to make organophosphorus
compounds, through the intermediates phosphorus
chlorides and two phosphorus sulfides, phosphorus
pentasulfide, and phosphorus sesquisulfide.
Organophosphorus compounds have many applications,
including in plasticizers, flame retardants,
pesticides, extraction agents, and water treatment.
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.
Matches
Phosphorus-contained matches were first produced
in the 1830s and contained a mixture of white
phosphorus, an oxygen-releasing compound (potassium
chlorate, lead dioxide or some nitrate) and
a binder in their heads. They were rather
sensitive to storage conditions, toxic and
unsafe, as they could be lit by striking on
any rough surface. Therefore, their production
was gradually banned between 1872 and 1925
in different countries. The international
Berne Convention, adopted in 1906, prohibited
the use of white phosphorus in matches. As
a consequence, the 'strike-anywhere' matches
were gradually replaced by 'safety matches'
where white phosphorus was replaced by phosphorus
sesquisulfide (P4S3), sulfur or antimony sulfide.
Such matches are hard to ignite on an arbitrary
surface and require a special strip. The strip
contains red phosphorus which heats up upon
striking, reacts with the oxygen-releasing
compound in the head and ignites the flammable
material of the head.
Water softening
Sodium tripolyphosphate made from phosphoric
acid is used in laundry detergents in some
countries, but banned for this use in others.
It is useful for softening water to enhance
the performance of the detergents and to prevent
pipe/boiler tube corrosion.
Niche applications
Phosphates are utilized in the making of special
glasses that are used for sodium lamps.
Bone-ash, calcium phosphate, is used in the
production of fine china.
Phosphoric acid made from elemental phosphorus
is used in food applications such as some
soft drinks. The acid is also a starting point
to make food grade phosphates. These include
mono-calcium phosphate that is employed in
baking powder and sodium tripolyphosphate
and other sodium phosphates. Among other uses
these are used to improve the characteristics
of processed meat and cheese. Others are used
in toothpaste.
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 signalling,
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 semiconductors.
32P and 33P are used as radioactive tracers
in biochemical laboratories.
Phosphate is a strong complexing agent for
the hexavalent uranyl (UO22+) species and
this is the reason why apatite and other natural
phosphates can often be very rich in uranium.
Tributylphosphate is an organophosphate soluble
in kerosene and used to extract uranium in
the Purex process applied in the reprocessing
of spent nuclear fuel.
Biological role
Inorganic phosphorus in the form of the phosphate
PO3−
4 is required for all known forms of life,
playing a major role in biological molecules
such as DNA and RNA where it forms part of
the structural framework of these molecules.
Living cells also use phosphate to transport
cellular energy in the form of adenosine triphosphate
(ATP). Nearly every cellular process that
uses energy obtains it in the form of ATP.
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.
Living cells are defined by a membrane that
separates it from its surroundings. Biological
membranes are made from a phospholipid matrix
and proteins, typically in the form of a bilayer.
Phospholipids are derived from glycerol, such
that two of the glycerol hydroxyl (OH) protons
have been replaced with 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% of which is
present 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. A well-fed adult
in the industrialized world 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
2PO−
4 and HPO2−
4. Only about 0.1% of body phosphate circulates
in the blood, and this amount reflects 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
fluoroapatite:
Phosphorus deficiency
In medicine, low-phosphate syndromes are caused
by malnutrition, by failure to absorb phosphate,
and by metabolic syndromes that draw phosphate
from the blood (such as re-feeding after malnutrition)
or pass too much of it into the urine. All
are characterized by hypophosphatemia, which
is a condition of low levels of soluble phosphate
levels in the blood serum, and therefore inside
cells. Symptoms of hypophosphatemia include
muscle and 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
in order to understand plant uptake from soil
systems. In ecological terms, phosphorus is
often a limiting factor in many environments;
i.e. the availability of phosphorus governs
the rate of growth of many organisms. In ecosystems
an excess of phosphorus can be problematic,
especially in aquatic systems, resulting in
eutrophication which sometimes lead to algal
blooms.
Food sources
The main food sources for phosphorus are foods
containing protein, although proteins do not
contain phosphorus. For example, milk, meat,
and soya typically also have phosphorus. As
a rule, if one holds a meal plan providing
sufficient amount of protein and calcium then
the amount of phosphorus is also likely sufficient.
Precautions
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 to certain organisms
as pesticides (herbicides, insecticides, fungicides,
etc.) and weaponised as nerve agents. 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". Ingestion
of white phosphorus may cause a medical condition
known as "Smoking Stool Syndrome".
Upon exposure to elemental phosphorus, in
the past it was suggested to wash the affected
area with 2% copper sulfate solution to form
harmless compounds that can be 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."
The manual suggests instead "a bicarbonate
solution to neutralize 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." 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.
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 as
List I precursor chemicals under 21 CFR 1310.02
effective on November 17, 2001. As a result,
in the United States, handlers of red or white
phosphorus are subject to stringent regulatory
controls.
