||silicon ← phosphorus → sulfur|
Periodic table - Extended periodic
|Name, symbol, number
||phosphorus, P, 15|
|Group, period, block
||15, 3, p|
||waxy white/ red/|
|Standard atomic weight
||[Ne] 3s2 3p3|
|Electrons per shell
||2, 8, 5|
|Density (near r.t.)
||(white) 1.823 g·cm−3|
|Density (near r.t.)
||(red) 2.34 g·cm−3|
|Density (near r.t.)
||(black) 2.69 g·cm−3|
||(white) 317.3 K|
(44.2 °C, 111.6 °F)
(277 °C, 531 °F)
|Heat of fusion
||(white) 0.66 kJ·mol−1|
|Heat of vaporization
||(25 °C) (white)|
Vapor pressure (white)
Vapor pressure (red)
||5, 4, 3, 2 , 1 , -3|
(mildly acidic oxide)
||2.19 (Pauling scale)|
|1st: 1011.8 kJ·mol−1|
|2nd: 1907 kJ·mol−1|
|3rd: 2914.1 kJ·mol−1|
|Atomic radius (calc.)
|Van der Waals radius
||(300 K) (white)|
|CAS registry number
Main article: Isotopes of phosphorus
||P is stable with 16 neutrons|
Phosphorus, (IPA: /ˈfɒsfərəs/, Greek: phôs meaning "light", and phoros meaning
"bearer"), is the chemical element that has the symbol P and atomic number 15. A multivalent nonmetal of the nitrogen group, phosphorus is commonly
found in inorganic phosphate rocks.
Due to its high reactivity, phosphorus is never found as
a free element in nature. One form of phosphorus (white phosphorus) emits a
faint glow upon exposure to oxygen (hence its Greek derivation and the Latin 'light-bearer', meaning the
planet Venus as
Hesperus or "Morning
Phosphorus is a component of DNA and RNA and an essential element for all
living cells. The most
important commercial use of phosphorus-based chemicals is the production of
Phosphorus compounds are also widely used in explosives, nerve agents, friction matches, fireworks, pesticides, toothpaste, and detergents.
Elemental phosphorus can exist in several allotropes, most commonly white, red
(P4) exists as individual molecules made up of four atoms in a
resulting in very high ring strain and instability. It contains 6 single bonds.
White phosphorus is a yellow, waxy transparent solid. For
this reason it is also called yellow phosphorus. It glows greenish in the dark
(when exposed to oxygen), is highly flammable and pyrophoric (self-igniting) upon contact
with air as well as toxic
(causing severe liver damage on ingestion). 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 P4O10
tetrahedra with oxygen inserted between the phosphorus atoms and at their
vertices. White phosphorus is insoluble in water but soluble in carbon
The white allotrope can be produced using several
different methods. In one process, calcium phosphate, which is derived
from phosphate rock, is heated in an electric or fuel-fired furnace in the
presence of carbon and
phosphorus is then liberated as a vapour and can be collected under phosphoric acid. This process is
similar to the first synthesis of phosphorus from calcium phosphate in
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 exists as an amorphous network of atoms which reduces strain and gives greater
stability; further heating results in the red phosphorus becoming crystalline.
Red phosphorus does not catch fire in air at temperatures below 240°C, whereas
white phosphorus ignites at about 30°C.
In 1865 Hittorf discovered that when phosphorus was recrystallized from molten
lead, a red/purple form is
obtained. This purple form is sometimes known as "Hittorf's phosphorus." In
addition, a fibrous form exists with similar phosphorus cages. Below is shown a
chain of phosphorus atoms which exhibits both the purple and fibrous
One of the forms of red/black phosphorus is a
Black phosphorus has an orthorhombic structure
(Cmca) and is the least reactive allotrope, it consists of many
six-membered rings which are interlinked. Each atom is bonded to three other
atoms. A recent synthesis of black phosphorus using metal salts as
catalysts has been reported.
The glow from phosphorus was the attraction of its
discovery around 1669, but the mechanism for that glow was not fully described
until 1974. It was known from early times that the glow 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 where 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 P2O2
that both emit visible light. The reaction is slow and only very little of the
intermediates is required to produce the luminescence, hence the extended time
the glow continues in a stoppered jar.
Although the term phosphorescence is derived from
phosphorus, the reaction which gives phosphorus its glow is properly called
luminescence (glowing by its own reaction, in this case chemoluminescence), not phosphorescence
(re-emitting light that previously fell on it).
Concentrated phosphoric acids, which can consist of 70%
to 75% P2O5 are very important to agriculture and farm production in the
form of fertilisers. Global demand for fertilizers led to large increases in
(PO43-) production in the second half of the 20th century.
- Phosphates are utilized in the making of special
glasses that are used for
- Bone-ash, calcium
phosphate, is used in the production of fine china.
- Sodium tripolyphosphate made
from phosphoric acid is used in laundry detergents in several countries, and
banned for this use in others.
- Phosphoric acid made from elemental phosphorus is used
in food applications such as soda beverages. The acid is also a starting point
to make food grade phosphates. These include mono-calcium phosphate which 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. Trisodium phosphate is used in cleaning agents to soften
water and for preventing pipe/boiler tube
- Phosphorus is widely used to make organophosphorus compounds, through the
intermediates phosphorus chlorides and the two phosphorus sulfides: phosphorus pentasulfide, and
Organophosphorus compounds have many applications, including in plasticizers, flame retardants, pesticides, extraction 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.
- White phosphorus is used in
military applications as
incendiary bombs, for
smoke-screening as smoke
pots and smoke bombs, and
in tracer ammunition.
- Red phosphorus is essential for manufacturing matchbook
strikers, flares, safety matches, pharmaceutical grade and street methamphetamine, and is used in
cap gun caps.
- Phosphorus sesquisulfide is used in heads of
- In trace amounts, phosphorus is used as a dopant for N-type semiconductors.
- 32P and 33P are used as
radioactive tracers in biochemical laboratories.
Phosphorus is a key element in all known forms of
life. Inorganic phosphorus
in the form of the phosphate PO43- plays 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 via 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.
An average adult human contains a little less than 1 kg
of phosphorus, about 85% of which is present in bones and teeth in the form of
apatite, and the remainder
inside cells in soft tissues. A well-fed adult in the industrialized world
consumes and excretes about 1-3 g of phosphorus per day in the form of
phosphate. Only about 0.1% of body phosphate circulates in the blood, but this
amount reflects the amount of phosphate available to soft tissue
In medicine, low phosphate syndromes are caused by
malnutrition, by failure to absorb phosphate, and by metabolic syndromes which
draw phosphate from the blood or pass too much of it into the urine. All are
characterized by hypophosphatemia (see article for medical details). Symptoms of low phosphate
include muscle and neurological dysfunction, and disruption of muscle and blood
cells due to lack of ATP.
Phosphorus is an essential macromineral for plants, which is
studied extensively in soil conservation in order to understand plant uptake from soil systems. In
phosphorus is often a limiting nutrient 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, see eutrophication and algal blooms.
Phosphorus (Greek phosphoros was the ancient name for the planet
Venus, but in Greek mythology, Hesperus and Eosphorus
could be confused with Phosphorus) was discovered by German alchemist Hennig Brand in 1669 through a
preparation from urine,
which contains considerable quantities of dissolved phosphates from normal
metabolism. Working in Hamburg, Brand attempted to distill some salts by evaporating urine, and in the
process produced a white material that glowed in the dark and burned
brilliantly. Since that time, phosphorescence has been used to describe
substances that shine in the dark without burning.
Phosphorus was first made commercially, for the match
industry, in the 19th century, by distilling off phosphorus vapor from
precipitated phosphates heated in a retort. The precipitated phosphates
were made from ground-up bones that had been de-greased and treated with strong
acids. This process became obsolete in the late 1890s when the electric arc furnace was adapted to
reduce phosphate rock.
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 giving off
luminous steam). In addition, exposure to the vapours gave match workers a
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 a Berne Convention,
requiring its adoption as a safer alternative for match manufacture.
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
Britain (hydrogen being
highly inflammable if it
can be ignited). During World War II, Molotov cocktails of benzene
and phosphorus 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 (see
precautions below). People
covered in it have been known to commit suicide due to the torment.
Today phosphorus production is larger than ever. It is
used as a precursor for various chemicals, in particular the herbicide
glyphosate sold under the
brand name Roundup.
Production of white phosphorus takes place at large facilities and it is
transported heated in liquid form. Some major accidents have occurred during
transportation, train derailments at Brownston,
Nebraska and Miamisburg,
Ohio led to large fires. The worst accident in recent
times was an environmental one in 1968 when phosphorus spilled into the sea from
a plant at Placentia Bay, Newfoundland.
Due to its reactivity with air and many other
oxygen-containing substances, phosphorus is not found free in nature but it is
widely distributed in many different minerals.
Phosphate rock, which is partially made of apatite (an
impure tri-calcium phosphate mineral), is an important commercial source of this
element. Large deposits of apatite are located in China, Russia, Morocco, Florida, Idaho, Tennessee, Utah, and elsewhere. Albright and Wilson in the United
Kingdom and their Niagara Falls plant, for instance, were using phosphate rock in the 1890s and
1900s from Connetable,
Tennessee and Florida; 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 affected by phosphate rock sales by China
and the entry of their long standing Moroccan phosphate suppliers into the
purified wet phosphoric acid business.
Organic compounds of phosphorus form a wide class of
materials, some of which 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 weaponized as nerve agents. Most
inorganic phosphates are relatively nontoxic and essential nutrients. For
environmentally adverse effects of phosphates see eutrophication and algal blooms.
The white phosphorus allotrope should be kept under water
at all times as it presents a significant fire hazard due to its extreme
reactivity with atmospheric oxygen, and it should only be manipulated with
forceps since contact with skin can cause severe burns. 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".
When the white form is exposed to sunlight or when it is
heated in its own vapour to 250°C, it is transmuted to the red form, which does
not phosphoresce in air. The red allotrope does not spontaneously ignite in air
and is not as dangerous as the white form. Nevertheless, it should be handled
with care because it reverts to white phosphorus in some temperature ranges and
it also emits highly toxic
fumes that consist of phosphorus oxides when it is heated.
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
The manual suggests instead "a bicarbonate solution to
neutralize phosphoric acid, which will then allow removal of visible WP.
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." Then, "Promptly debride the burn if the patient's
condition will permit removal of bits of WP which 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.
Further warnings of toxic effects and recommendations for
treatment can be found in the Emergency War Surgery NATO Handbook: Part I:
Types of Wounds and Injuries: Chapter III: Burn Injury: Chemical Burns And White
DEA List I
Phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent
effective for reducing ephedrine or pseudoephedrine to methamphetamine. For this reason, two allotropes of elemental
phosphorus—red phosphorus and white phosphorus—were designated by the United
States Drug Enforcement Administration as List I precursor
chemicals under 21 CFR
1310.02 effective November 17, 2001. As a result, in
the United States,
handlers of red phosphorus or white phosphorus are subject to stringent
regulatory controls pursuant to the Controlled
Substances Act in order to reduce diversion of these
substances for use in clandestine production of controlled
As an exception to the octet
The simple Lewis
structure for the trigonal
bipyramidal PCl5 molecule contains five
covalent bonds, implying a
hypervalent molecule with
ten valence electrons contrary to the octet
An alternate description of the bonding, however,
respects the octet rule by using 3-center-4-electron
(3c-4e) bonds. In this model the octet on the P atom
corresponds to six electrons which form three Lewis (2c-2e) bonds to the three
equatorial Cl atoms, plus the two electrons in the 3-centre Cl-P-Cl bonding
molecular orbital for the two axial Cl electrons. The two electrons in the
corresponding nonbonding molecular orbital are not included because this orbital
is localized on the two Cl atoms and does not contribute to the electron density on P.
Radioactive isotopes of phosphorus
- 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
Occupational Safety and Health
Administration requires 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 a process
known as Bremsstrahlung,
meaning braking radiation.
Therefore shielding must be accomplished with low density materials, e.g.
Plexiglas, Lucite, plastic, wood, or water.
- 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.
According to the Oxford English Dictionary the correct
spelling of the element is phosphorus. The word phosphorous is the
adjectival form for the P3+ valency: so, just as sulfur forms sulfurous and
sulfuric compounds, phosphorus forms phosphorous and
- Hydride: PH3
- Halides: PBr5, PBr3, PCl3, PI3
- Oxides: P4O10
- Sulfides: P2S5,
- Acids: H3PO2,
- Phosphates: (NH4)3PO4, Ca3(PO4)2), FePO4, Fe3(PO4)2, Na3PO4,
- Phosphides: Ca3P2,
- Organophosphorus and
Parathion, Sarin, Soman, Tabun, Triphenyl phosphine, VX nerve gas
By: Zookeeper - 2007-12-08 19:48:18