|Name, Symbol, Number|| phosphorus, P, 15
|Group, Period, Block||15, 3, p|
|Appearance|| waxy white/ red/|
|Atomic mass||30.973762(2) g/mol|
|Electron configuration||[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|
|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−1|
|Heat of vaporization||12.4 kJ·mol−1|
|Heat capacity|| (25 °C) (white)|
|Oxidation states|| ±3, 5, 4|
(mildly acidic oxide)
|Electronegativity||2.19 (Pauling scale)|
| Ionization energies|
|1st: 1011.8 kJ·mol−1|
|2nd: 1907 kJ·mol−1|
|3rd: 2914.1 kJ·mol−1|
|Atomic radius||100 pm|
|Atomic radius (calc.)||98 pm|
|Covalent radius||106 pm|
|Van der Waals radius||180 pm|
|Magnetic ordering||no data|
|Thermal conductivity|| (300 K) (white)|
|Bulk modulus||11 GPa|
|CAS registry number||7723-14-0|
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.
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.
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, too high or too low and the reaction stops. 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 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).
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 led to large increases in phosphate (PO43-) 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, also known as 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 elementary 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 corrosion.
- Phosphorus is widely used to make organophosphorus compounds, through the intermediates phosphorus chlorides and the two phosphorus sulfides: phosphorus pentasulfide, and phosphorus sesquisulfide. 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 and, most notoriously, methamphetamine.
- In trace amounts, phosphorus is used as a dopant for N-type semiconductors.
- 32P and 33P are used as radioactive tracers in biochemical laboratories (see Isotopes).
- Red phosphorus is used in cap gun caps.
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 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.
Phosphorus (Greek name = phosphoros, was the ancient name for the planet Venus) was discovered by German alchemist Henning 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 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 electriv 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 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. 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 of course being highly flammable 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 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.
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. 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.
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.
See also Phosphate minerals.
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, insectinides, 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 allotrope white phosphorus should be kept under water at all times as it presents a significant fire hazard 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". 
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.
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) 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 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 disrecommended 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 Phosphorus injury.
- For more details on this topic, see 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 nucleic acids, OSHA 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, acrylic, 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, so phosphorus forms phosphorous and phosphoric 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
- 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)
- Triphenyl phosphine
- Monopotassium phosphate (KH2PO4)
- Trisodium phosphate (Na3PO4)
- VX nerve gas
See also Phosphorus compounds.
- ↑ 1.0 1.1 1.2 Emsley, John (2000). The Shocking History of Phosphorus, ISBN 0-330-39005-8
- ↑ Nobel Prize in Chemistry 1956 - Presentation Speech, by Professor A. Ölander (committee member)
- ↑ Phosphorus Topics page, at Lateral Science
- ↑ 4.00 4.01 4.02 4.03 4.04 4.05 4.06 4.07 4.08 4.09 4.10 4.11 4.12 4.13 Threlfall, R.E., (1951). 100 years of Phosphorus Making: 1851 - 1951. Oldbury: Albright and Wilson Ltd
- ↑ Aall C. H. (1952). "The American Phosphorus Industry". Industrial & Engineering Chemistry 44. DOI:10.1021/ie50511a018.
- ↑ Podger, Hugh, (2002). Albright & Wilson: The Last 50 Years. Studley: Brewin Books. ISBN 1-85858-223-7
- ↑ emedicine.com CBRNE - Incendiary Agents, White Phosphorus (Smoking Stool Syndrome)
- ↑ US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries
- ↑ Emergency War Surgery NATO Handbook: Part I: Types of Wounds and Injuries: Chapter III: Burn Injury: Chemical Burns And White Phosphorus injury.
- ↑ http://www.oseh.umich.edu/TrainP32.pdf
- 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.
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