Phosphorus

Silicon - Phosphorus - Sulfur N P As       Full table
General
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 (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 ×10-3 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
Miscellaneous
Electronegativity	2.19 (Pauling scale)
Specific heat capacity	769 J/(kg*K)
Electrical conductivity	1.0 10-9/m ohm
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 t½ is 25.34 d (P-32) 31P 100% P is stable with 16 neutrons
SI units & STP are used except where noted.

Phosphorus is a 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 but is never naturally found alone. It is highly reactive, gives-off a faint glow upon uniting with oxygen (hence its name), occurs in several forms and is an essential element for living organisms. The most important use of phosphorus is in the production of fertilizers. It is also widely used in explosives, friction matches, fireworks, pesticides, toothpaste and detergents.

Table of contents 1 Notable Characteristics 1.1 Forms 2 Applications 3 Biological Role 4 History 5 Occurrence 6 Precautions 7 External Links

Notable Characteristics Common phosphorus forms a waxy white solid that has a characteristic disagreeable smell but when it is pure it is 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.

Forms

Phosphorus exists in four or more allotropic forms: white (or yellow), red, and black (or violet). The most common are red and white phosphorus, both of which are tetrahedral groups of four atoms. White phosphorus burns on contact with air and on exposure to heat or light it can transform into red phosphorus. It also exists in two modifications: alpha and beta which are separated by a transition temperature of -3.8 °C. Red phosphorus is comparatively stable and sublimes at a vapor pressure of 1 atm at 17 °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. Applications 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 production in the second half of the 20th century. Other uses;

• 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[?] and to make mono-calcium phosphate which is employed in baking powder.
• This element is also an important component in steel production, in the making of phosphor bronze, and in many other related products.
• Trisodium phosphate[?] is widely used in cleaning agents to soften water[?] and for preventing pipe/boiler tube corrosion.
• White phosphorus is used in military incendiaries, smoke pots, smoke bombs and tracer bullets.
• Miscellaneous uses; used in the making of safety matches, pyrotechnics, pesticides, toothpaste, detergentsetc.

Biological Role Phosphorus compounds perform vital functions in all known forms of life. Inorganic phosphorus plays a key role in biological molecules such as DNA and RNA where it forms part of those molecules' molecular backbones. Living cells also utilize inorganic phosphorus to store and transport cellular energy via adenosine triphosphate (ATP). Calcium phosphate salts are used by animals to stiffen bones and phosphorus is also an important element in cell protoplasm and nervous tissue. History 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. Occurrence 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. Precautions This is a particularly poisonous element with 50 mg[?] being the average fatal dose. The allotrope white phosphorus should be kept under water at all times due to its hyper reactivity to air 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[?]". Phosphate esters are nerve poisons but 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.

External Links

• WebElements.com - Phosphorus (http://www.webelements.com/webelements/elements/text/P/index)
• EnvironmentalChemistry.com - Phosphorus (http://environmentalchemistry.com/yogi/periodic/P)