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15 silicon ← phosphorus → sulfur N ↑ P ↓ As Image:P-TableImage.png periodic table
General
Name, Symbol, Number	phosphorus, P, 15
Chemical series	nonmetals
Group, Period, Block	15, 3, p
Appearance	waxy white/ red/ black/ colorless Image:P,15.jpg
Atomic mass	30.973762(2) g/mol
Electron configuration	[Ne] 3s2 3p3
Electrons per shell	2, 8, 5
Physical properties
Phase	solid
Density (near r.t.)	(white) 1.823 g/cm³
Density (near r.t.)	(red) 2.34 g/cm³
Density (near r.t.)	(black) 2.69 g/cm³
Melting point	(white) 317.3 K (44.2 °C, 111.6 °F)
Boiling point	550 K (277 °C, 531 °F)
Critical temperature	994 K
Heat of fusion	(white) 0.66 kJ/mol
Heat of vaporization	12.4 kJ/mol
Heat capacity	(25 °C) (white) 23.824 J/(mol·K)
Vapor pressure (white) P/Pa 1 10 100 1 k 10 k 100 k at T/K 279 307 342 388 453 549
Vapor pressure (red) P/Pa 1 10 100 1 k 10 k 100 k at T/K 455 489 529 576 635 704
Atomic properties
Oxidation states	±3, 5, 4 (mildly acidic oxide)
Electronegativity	2.19 (Pauling scale)
Ionization energies (more)	1st: 1011.8 kJ/mol
	2nd: 1907 kJ/mol
	3rd: 2914.1 kJ/mol
Atomic radius	100 pm
Atomic radius (calc.)	98 pm
Covalent radius	106 pm
Van der Waals radius	180 pm
Miscellaneous
Magnetic ordering	no data
Thermal conductivity	(300 K) (white) 0.236 W/(m·K)
Bulk modulus	11 GPa
CAS registry number	7723-14-0
Notable isotopes
Main article: [[Isotopes of {{{isotopesof}}}]] iso NA half-life DM DE (MeV) DP 31P 100% P is stable with 16 neutrons 32P syn 14.28 d β- 1.709 32S 33P syn 25.3 d β- 0.249 33S
References

Phosphorus, (IPA: /ˈfɒsfərəs/, Greek: phôs

meaning "light", and phoros meaning "bearer"), 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. 

Phosphorus exists in several allotropes, most commonly white, red and black. White phosphorus (P₄) contains only four atoms, resulting in very high ring strain and instability. White phosphorus glows in the dark, is highly flammable and pyrophoric (self-igniting) upon contact with air as well as toxic. Red phosphorus has a network form which reduces strain and gives greater stability. Red phosphorus does not catch fire in air at temperatures below 240°C whereas white phosphorus ignites at about 40°C. Black phosphorus is amorphous and is the least reactive allotrope.

Red phosphorus is formed by heating white phosphorus to 250°C (482°F) or by exposing white phosphorus to sunlight.

Due to its high reactivity, phosphorus is never found as a free element in nature. It emits a faint glow upon exposure to oxygen (hence its Greek derivation and the Latin meaning 'morning star') and is an essential element for living organisms. The most important commercial use of phosphorus-based chemicals is the production of fertilizers. They are also widely used in explosives, nerve agents, friction matches, fireworks, pesticides, toothpaste, and detergents.

Contents 1 Characteristics 1.1 Glow 2 Biological role 3 History 4 Occurrence 5 Spelling 6 Compounds 7 References 8 External links

[edit] Characteristics

Phosphorus, in its common form, is a waxy white (or yellowish) solid that has a characteristic, disagreeable smell similar to that of garlic. Pure forms of the element are colorless and transparent. This nonmetal is not soluble in water, but is soluble in carbon disulfide. The white allotrope ignites spontaneously in air; however both white and red phosphorus burn in air to produce phosphorus pentoxide.

[edit] Glow

The glow from phosphorus was the attraction of its discovery around 1669, but the mechanism for that glow was not fully described until 1974.^[1] 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,^[2] there is only a range of partial pressure where it does, too high or too low and the reaction stops. Heat can be applied to drive the reaction at higher pressures.^[3]

In 1974 the glow was explained by R. J. van Zee and A. U. Khan.^[1] A reaction with oxygen takes place at the surface of the solid (or liquid) phosphorus, forming short-lived molecules HPO and P₂O₂ and they 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 is properly called luminescence (glowing by its own reaction, in this case chemoluminescence), not phosphorescence (re-emitting light that previously fell on it).

[edit] Biological role

Phosphorus is a key element in all known forms of life. Inorganic phosphorus in the form of the phosphate PO₄^3- 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 utilize phosphate to transport cellular energy via adenosine triphosphate (ATP). Nearly every cellular process that uses energy gets 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 are used by animals to stiffen their bones. An average person contains a little less than 1 kg of phosphorus, about three quarters of which is present in bones and teeth in the form of apatite. A well-fed adult in the industrialized world consumes and excretes about 1-3 g of phosphorus per day in the form of phosphate. Phosphorus is an essential mineral macronutrient, which is studied extensively in soil conservation in order to understand plant uptake from soil systems.

In ecological terms, 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.

[edit] History

Phosphorus (Greek phosphoros 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.

Phosphorus was first made commercially, for the match industry, in the 19th century, by distilling off phosphorus vapour from precipitated phosphates heated in a retort^[4] The precipitated phosphates were made from ground-up bones that had been de-greased and treated with strong acids^[4]. This process became obsolete in the late 1890s when the electric arc furnace was adapted to reduce phosphate rock^[4].

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 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 phosphorus could be used in weapons of war.^[1][4] In World War I it was used in incendiaries, smoke screens and tracer bullets^[4]. A special incendiary bullet was developed to shoot at hydrogen filled Zeppelins over Britain (hydrogen of course being highly flammable if it can be ignited)^[4]. 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 were known to commit suicide due to the torment.

Today phosphorus production is larger than ever, 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 is transported heated in liquid form. Some major accidents have occurred during transportation, train derailments at Brownston, Nebraska and Miamisburg, Ohio lead to large fires. The worst accident in recent times though was an environmental one in 1968 when phosphorus spilt into the sea from a plant at Placentia Bay, Newfoundland.

[edit] 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 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; however, by 1950 they were using phosphate rock mainly from Tennessee and North Africa^[4]. 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^[5].

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^[4]. Elemental phosphorus is then liberated as a vapor and can be collected under phosphoric acid.

See also Phosphate minerals.

[edit] Spelling

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, so phosphorus forms phosphorous and phosphoric compounds.

[edit] Compounds

• Ammonium phosphate ((NH₄)₃PO₄)
• Calcium phosphate (Ca₃(PO₄)₂)
• Calcium dihydrogen phosphate (Ca(H₂PO₄)₂)
• Calcium phosphide (Ca₃P₂)
• Iron(III) phosphate (FePO₄)
• Iron(II) phosphate (Fe₃(PO₄)₂)
• Gallium(III) phosphide (GaP)
• Hypophosphorous acid (H₃PO₂)
• Lawesson's reagent
• Parathion
• Phosphine (Phosphorus Trihydride PH₃)
• Phosphoric acid (H₃PO₄)
• Phosphorus pentabromide (PBr₅)
• Phosphorus pentasulfide (P₂S₅)
• Phosphorus pentoxide (P₂O₅)
• Phosphorus sesquisulfide (P₄S₃)
• Phosphorus tribromide (PBr₃)
• Phosphorus trichloride (PCl₃)
• Phosphorus triiodide (PI₃)
• Sarin
• Soman
• Tabun
• Triphenyl phosphine
• Monopotassium phosphate (KH₂PO₄)
• Trisodium phosphate (Na₃PO₄)
• VX nerve gas

See also Phosphorus compounds.

[edit] References

• Los Alamos National Laboratory – Phosphorus

1. ↑ ^{1.0 1.1 1.2} Emsley, John (2000). The Shocking History of Phosphorus, ISBN 0-330-39005-8
2. ↑ Nobel Prize in Chemistry 1956 - Presentation Speech, by Professor A. Ölander (committee member)
3. ↑ Phosphorus Topics page, at Lateral Science
4. ↑ Cite error: Invalid <ref> tag; no text was provided for refs named threlfall
5. ↑ Podger, Hugh, (2002). Albright & Wilson: The Last 50 Years. Studley: Brewin Books. ISBN 1-85858-223-7

[edit] External links

• WebElements.com – Phosphorus
• Entrez PubMed - Acute Yellow Phosphorus Poisoning
• eMedicine.com - Article on White Phophorus as used as weapon
• Link page to external chemical sources.af:Fosfor

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