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This article is about the chemical element. For the article about Phosphorus meaning "morning star", go to Phosphorus (morning star).

Name, Symbol, Number phosphorus, P, 15
Chemical series Nonmetals
Group, Period, Block 15 (VA), 3, p
Density, Hardness 1823 kg/m3, __
Appearance colorless/red/silvery white
Atomic properties
Atomic weight 30.973761 amu
Atomic radius (calc.) 100 pm (98 pm)
Covalent radius 106 pm
van der Waals radius 180 pm
Electron configuration [Ne]3s2 3p3
e- 's per energy level 2, 8, 5
Oxidation states (Oxide) 3, 5, 4 (mildly acidic)
Crystal structure monoclinic
Physical properties
State of matter Solid
Melting point 317.3 K (111.6 F)
Boiling point 550 K (531 F)
Molar volume 17.02 ×10-6 m3/mol
Heat of vaporization 12.129 kJ/mol
Heat of fusion 0.657 kJ/mol
Vapor pressure 20.8 Pa at 294 K
Speed of sound no data
Electronegativity 2.19 (Pauling scale)
Specific heat capacity 769 J/(kg*K)
Electrical conductivity 1.0 10-9/(mohm)
Thermal conductivity 0.235 W/(m*K)
1st ionization potential 1011.8 kJ/mol
2nd ionization potential 1907 kJ/mol
3rd ionization potential 2914.1 kJ/mol
4th ionization potential 4963.6 kJ/mol
5th ionization potential 6273.9 kJ/mol
Most stable isotopes
iso NA Longest t is 25.34 d (P-33)
31P 100% P is stable with 16 neutrons
SI units & STP are used except where noted.

Phosphorus is the chemical element in the periodic table that has the symbol P and atomic number 15. A multivalent, nonmetal of the nitrogen group, phosphorus is commonly found in inorganic phosphate rocks and in all living cells. Due to its high reactivity, it is never found as a free element in nature. It emits a faint glow upon exposure to oxygen (hence its name, Latin for 'morning star', from Greek words meaning 'light' and 'bring'), occurs in several allotropic forms, and is an essential element for living organisms. The most important commercial use of phosphorus is in the production of fertilizers. It is also widely used in explosives, friction matches, fireworks, pesticides, toothpaste, and detergents.


Notable characteristics

Common phosphorus forms a waxy white solid that has a characteristic disagreeable smell. Pure forms of the element are colorless and transparent. This non metal is not soluble in water, but it is soluble in carbon disulfide. Pure phosphorus ignites spontaneously in air and burns to phosphorus pentoxide.


Phosphorus exists in four allotropic forms: white (or yellow), red, and black (or violet). Other allotropic forms may exist. The most common are red and white phosphorus, both of which consist of networks of tetrahedrally arranged groups of four phosphorus atoms. The tetrahedra of white phosphorous form separate groups; the tetrahedra of red phosphorus are linked into chains. White phosphorus burns on contact with air and, on exposure to heat or light, it can transform into red phosphorus.

Phosphorus also exists in kinetically and thermodynamically favored forms. They are separated by a transition temperature of -3.8 C. One is known as the "alpha" form, the other "beta". Red phosphorus is comparatively stable and sublimes at a vapor pressure of 1 atm at 170 C but burns from impact or frictional heating. A black phosphorus allotrope exists which has a structure similar to graphite - the atoms are arranged in hexagonal sheet layers and will conduct electricity.


Concentrated phosphoric acids, which can consist of 70% to 75% P2O5 are very important to agriculture and farm production in the form of fertilizers. Global demand for fertilizers has led to large increases in phosphate (PO43-) production in the second half of the 20th century. Other uses;

Biological role

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 backbone of these molecules. Living cells also utilize phosphate to transport cellular energy via adenosine triphosphate (ATP). Nearly every cellular process that uses energy gets it in the form of ATP. Phospholipids are the main structural components of all cellular membranes. Calcium phosphate salts are used by animals to stiffen their bones.


Phosphorus (Greek. phosphoros, meaning "light bearer" which was the ancient name for the planet Venus) was discovered by German alchemist Hennig Brand in 1669 through a preparation from urine. Working in Hamburg, Brand attempted to distill 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.

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 vapors gave match workers a necrosis of the bones of the jaw, the infamous "phossy-jaw." When red phosphorus was discovered, with its far lower flammability and toxicity, it was adopted as a safer alternative for match manufacture.


Due to its reactivity to 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 in Russia, Morocco, Florida, Idaho, Tennessee, Utah, and elsewhere.

The white allotrope can be produced using several different methods. In one process, tri-calcium phosphate, which is derived from phosphate rock, is heated in an electric or fuel-fired furnace in the presence of carbon and silica. Elemental phosphorus is then liberated as a vapor and can be collected under phosphoric acid.


This is a particularly poisonous element with 50 mg being the average fatal dose (white phosphorus is generally considered to be the lethal form of phosphorus while phosphate and orthophosphate are essential nutrients). The allotrope white phosphorus should be kept under water at all times due to its extreme reactivity to atmospheric oxygen, and it should only be manipulated with forceps since contact with skin can cause severe burns. Chronic white phosphorus poisoning of unprotected workers leads to necrosis of the jaw called "phossy-jaw". Ingestion of white phosphorus may cause a medical condition known as "Smoking Stool Syndrome". Fluorophosphate esters are among the most potent neurotoxins known but most inorganic phosphates are relatively nontoxic. Phosphate pollution occurs where fertilizers or detergents have leached into soils.

When the white form is exposed to sunlight or when it is heated in its own vapor 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 does revert to white phosphorus in some temperature ranges and it also emits highly toxic fumes that consist of phosphorus oxides when it is heated.


Some common isotopes of phosphorus include:

  • 32P (radioactive). Phosphorus-32 is a beta-emitter (1.71 MeV) with a half-life of 14.3 days. It is used routinely in life-science laboratories, primarily to produce radiolabeled DNA and RNA probes, typically for use in Northern blots or Southern blots.
  • 33P (radioactive). Phosphorus-33 is 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, for example, DNA sequencing.


The only correct spelling of the element is phosphorus. There does exist a word phosphorous, but it is the adjectival form for the smaller valency: so, just as sulfur forms sulfurous and sulfuric compounds, so phosphorus forms phosphorous and phosphoric compounds.


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