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phosphorus

  (fŏs'fər-əs) pronunciation
n.
  1. (Symbol P) A highly reactive, poisonous, nonmetallic element occurring naturally in phosphates, especially apatite, and existing in three allotropic forms, white (or sometimes yellow), red, and black. An essential constituent of protoplasm, it is used in safety matches, pyrotechnics, incendiary shells, and fertilizers and to protect metal surfaces from corrosion. Atomic number 15; atomic weight 30.9738; melting point (white) 44.1°C; boiling point 280°C; specific gravity (white) 1.82; valence 3, 5.
  2. A phosphorescent substance.

[Latin Phōsphorus, morning star, from Greek phōsphoros, bringing light, morning star : phōs, light + -phoros, -phorous.]


 
 

A chemical element, P, atomic number 15, atomic weight 30.9738. Phosphorus forms the basis of a very large number of compounds, the most important class of which are the phosphates. For every form of life, phosphates play an essential role in all energy-transfer processes such as metabolism, photosynthesis, nerve function, and muscle action. The nucleic acids which among other things make up the hereditary material (the chromosomes) are phosphates, as are a number of coenzymes. Animal skeletons consist of a calcium phosphate. See also Periodic table.

About three-quarters of the total phosphorus (in all of its chemical forms) used in the United States goes into fertilizers. Other important uses are as builders for detergents, nutrient supplements for animal feeds, water softeners, additives for foods and pharmaceuticals, coating agents for metal-surface treatment, additives in metallurgy, plasticizers, insecticides, and additives for petroleum products.

Of the nearly 200 different phosphate minerals, only one, fluorapatite, Ca5F(PO4)3, is mined chiefly from large secondary deposits originating from the bones of dead creatures deposited on the bottom of prehistoric seas and from bird droppings on ancient rookeries.

Research in phosphorus chemistry indicates that there may be as many compounds based on phosphorus as on carbon. In organic chemistry it has been customary to group the various chemical compounds based on carbon into families which are called homologous series. This can also be done in the chemistry of phosphorus compounds, even though many phosphorus-based families are incomplete. The best known of the families of compounds based on phosphorus is the group of chain phosphates. Phosphate salts consist of cations, such as sodium, along with chain anions, such as (PnO3n+1)(n−2)−, which may have 1–1,000,000 phosphorus atoms per anion.

The phosphates are based on phosphorus atoms tetrahedrally surrounded by oxygen atoms, with the lowest member of the series being the simple PO43− anion (the orthophosphate ion). The family of chain phosphates is based on a row of alternating phosphorus and oxygen atoms in which each phosphorus atom remains in the center of a tetrahedron of four oxygen atoms. There is also a closely related family of ring phosphates, a member of which, the trimetaphosphate, is shown in Fig. 1.

Ring phosphate anion, (P<sub>3</sub>O<sub>9</sub>)<SUP>3−</SUP>.
Ring phosphate anion, (P3O9)3−.

An interesting structural characteristic of many known phosphorus compounds is the formation of cagelike structures. Such cagelike molecules are exemplified by white phosphorus, P4, and one of the phosphorus pentoxides, P4O10 (Fig. 2). Network structures are also common; for example, black phosphorus crystals in which the atoms are bonded together in the form of vast, corrugated planes (Fig. 3).

Phosphorus pentoxide, P<sub>4</sub>O<sub>10</sub>, in vapor state.
Phosphorus pentoxide, P4O10, in vapor state.

Black phosphorus, P<sub><i>n</i></sub>.
Black phosphorus, Pn.

In the majority of its compounds, phosphorus is chemically bonded to four neighboring atoms. There is a large number of compounds in which one of the four neighboring atoms is absent, and in which its place is taken by an unshared pair of electrons. There are also a few compounds in which there are five or six neighboring atoms bonded to the phosphorus. These compounds are very reactive and tend to be unstable.

During the 1960s and 1970s a large number of organic-phosphorus compounds were prepared. Most of these chemical structures involve three or four neighboring atoms bonded to the phosphorus, but stable structures having two, five, or six neighboring atoms per phosphorus are also known. See also Organophosphorus compound.

Essentially all of the phosphorus used in commerce is in the form of phosphates. The majority of phosphatic fertilizers consist of highly impure monocalcium or dicalcium orthophosphate, Ca(H2PO4)2 and CaHPO4. These phosphates are salts of orthophosphoric acid.

The phosphorus compound of major biological importance is adenosine triphosphate (ATP), which is an ester of sodium tripolyphosphate, widely employed in detergents and water-softening compounds. Practically every reaction in metabolism and photosynthesis involves the hydrolysis of this tripolyphosphate to its pyrophosphate derivative, called adenosine diphosphate (ADP). See also Adenosine triphosphate (ATP).


 
Food and Nutrition: phosphorus

An essential element, occurring in tissues and foods as phosphate (salts of phosphoric acid), phospholipids, and phosphoproteins. In the body most (80%) is present in the skeleton and teeth as calcium phosphate (hydroxyapatite); the remainder is in the phospholipids of cell membranes, in nucleic acids, and in a variety of metabolic intermediates, including ATP. The parathyroid hormone controls the concentration of phosphate in the blood, mainly by modifying its excretion in the urine.

Adult needs (about 1.3 g per day) are always met. The calcium to phosphate ratio of infant foods is, however, important. Phosphate deficiency is common in livestock and gives rise to osteomalacia (also known as sweeny or creeping sickness).

 
Food and Fitness: phosphorus

A non-metallic element which is an essential component of the diet. It is a constituent of many vital compounds in the body, including ATP, DNA, and phospholipids (see separate entries), but it is found mainly in bones. Vitamin D and calcium regulate the availability of phosphorus for bone formation. Most meats and fish are rich sources of phosphorus; deficiencies lead to rickets and poor growth, but this is rare compared with the effects of calcium and vitamin D deficiency on bone development.

 
Dental Dictionary: phosphorus
(fos′fər-us)
n
P

A nonmetallic element; atomic weight, 30.98. It is essential, as is the phosphate, for the mineralization of the organic matrix of teeth and bone. It is also essential in the intermediary metabolism of carbohydrates as a vital constituent of the various intermediary compounds (e.g., glucose 6-phosphate) and of the enzyme systems (e.g., adenosine triphosphate [ATP]).

 

Description

Phosphorus (chemical symbol P) is a mineral discovered by the German alchemist Hennig Brand in 1669. It plays an essential part in multiple biochemical reactions for both plants and animals and is essential to all life. Phosphorus is found in living things, in soil and rock, mostly as chemical compounds known as phosphates. Rock and soil phosphorus are mined extensively throughout the world, but especially in the Peoples' Republic of China and the United States.

Phosphorus extracted from rock is classified as either white, red or black. White (also called yellow or common) phosphorus is a wax-like substance created by heating phosphate rock until it vaporizes and the condensation solidifies. One of this form's characteristics has given the English language the adjective phosphorescent, from white phosphorus's capacity to glow in the dark when exposed to air.

White phosphorus is highly toxic, causes burns if it comes in contact with skin, and is so combustible that it has to be stored underwater for safety. Red phosphorus is a rust-colored powder created by heating white phosphorus and exposing it to sunlight. It is not as combustible as the white form. Black phosphorus is made by heating white phosphorus under extremely high pressure until it resembles graphite.

In plants, phosphorus is necessary for photosynthesis to take place. In the human body, phosphorus works in tandem with another element, calcium, in much the same way that two other electrolyte components, sodium and potassium, do. Though phosphorus is found in every cell of the human body and accounts for 1% of the body's total weight, its primary function is working in conjunction with calcium to form teeth and bones.

Eighty-five percent of the phosphorus found in the body is located in these structures. In a delicately balanced chemical reaction, substances known as PTH (parathyroid hormone), Calcitonin, and 25-Dihydroxy vitamin D regulate the absorption of both calcium and phosphorus from the intestinal tract, thus making it available for the production of bones and teeth. If an excessive amount of phosphorus is absorbed, this will result in the phosphorus combining with all available calcium and preventing the calcium's efficient use in making and maintaining bones and teeth.

PTH balances the proportions of calcium and phosphorus in the body by increasing the release of calcium and phosphate from bone and the loss of phosphorus via the kidneys while limiting the excretion of calcium. PTH also increases the activity of the 25-Dihydroxy v25-Dihydroxy vitamin D, which, in contrast, increases the absorption of both phosphorus and calcium from the intestinal tract.

General Use

White phosphorus is a component of fertilizers, detergents and water softeners. It is also used in the manufacture of steel, plastics, insecticides, medical drugs, and animal feeds. Both white and red phosphorus are used in the making of safety matches and pesticides, including rat poison.

But the 15% of this element found in the blood stream and in other soft tissue also has a highly significant part to play in a variety of other body functions. Working with Vitamin B, phosphorus is involved in the metabolism of fats and carbohydrates, in both the repair of damaged cells and tissues and the routine maintenance of healthy ones. Phosphorus is necessary for the regularity of the heartbeat, and aids in the contraction of all other muscles throughout the body. Phosphorus is needed for the functioning of the kidneys and plays a part in the conduction of impulses along the network that makes up the nervous system.

Preparations

According to the American Dietetic Association, phosphorus intake in the United States is generally above what is needed, and in recent years has actually increased. Therefore, under normal circumstances with normal food intake, there is seldom if ever a need to supplement intake of phosphorus. Persons suffering from eating disorders such as anorexia and bulimia can be deficient in phosphorus intake as well as other nutrients. As the best source of phosphorus is in protein foods such as meat, eggs and milk products, some vegetarians may also need to evaluate their intake of this element. Excess consumption of processed foods, and inadequate intake of whole foods, plus fertilizers and pesticides are some of the causes for excess phosphorus.

Beside high-protein foods, phosphorus is also found in decreasing quantities in whole grain breads and cereals, especially unprocessed ones, and in minute amounts in fruits and vegetables. The phosphorus present in whole grain breads and cereals, however, exists as a substance called phytin. Phytin combines with calcium to create a salt that the human body is incapable of absorbing, thus making unprocessed, unenriched grains a negligible source of phosphorus. But both commercially prepared cereals and breads may provide this element as they are frequently enriched with it. Phosphates can also be taken by mouth as a tablet.

Precautions

White phosphorus is poisonous. Red phosphorus is not. As noted, white phosphorus is a highly toxic, flammable substance capable of burning the skin if it makes contact, and of igniting at room temperature. It should be handled with extreme care. Accidental phosphorus poisoning can happen from both fertilizers and pesticides. Phosphates sometimes are leached into water systems through sewage and can drastically alter the chemical makeup of lakes and rivers. In sufficient quantities, they can lead to the death of nearly all forms of aquatic life.

A normal blood serum level of phosphorus is 2.4-4.1 mg per deciliter of blood. An abnormal serum phosphorus level should be evaluated by a physician.

Phosphorus levels higher than normal can indicate a diet that includes an excessive phosphorus intake, inadequate intake of calcium, or lack of PTH (parathyroid hormone) in the system. It can be related to bone metastasis associated with cancer, liver or kidney disease, or sarcoidosis.

Serum phosphorus levels that are below normal can be related to insufficient phosphorus or vitamin D in one's diet leading to rickets in children and osteomalacia in adults. Disorders of the parathyroid gland, causing it to secrete excessive quantities of PTH, or of the pancreas, causing it to secrete too much insulin, also affect blood levels of phosphorus. Diabetic ketoacidosis or too much calcium are other possible causes. Multiple endocrine neoplasia (MEN) is yet another condition that often is associated with lower than normal levels of phosphorus.

Side Effects

Phosphorus preparations taken to supplement low phosphorus levels in the body can cause diarrhea.

Interactions

Antacids can decrease the absorption of phosphorus. Laxatives and enemas that contain the chemical compound sodium phosphate and excessive intake of vitamin D can increase phosphorus levels in the body. Administration of intravenous glucose solutions will cause phosphorus to combine with the glucose that is being absorbed by the cells.

Resources

Books

Busch, Marianna A., Ph.D. Phosphorus, World Book. Chicago, IL: World Book, Inc., 1999.

Clayman, Charles B., MD. The American Medical Association Home Medical Encyclopedia. New York: Random House, 1989.

Periodicals

Affenito, Sandra G., pH, RD, and Jane Kerstetter, pH, RD. "Position of the American Dietetic Association and Dietitians of Canada: Women's Health and Nutrition." Journal of the American Dietetic Association 1999.

Other

"Phosphorus in the Diet." http//www.Healthcentral.com. (June 2000).

"Serum Phosphorus." http//www.Healthcentral.com. (June 2000).

"Vitamin D." http//www.Healthcentral.com. (June 2000).

[Article by: Joan Schonbeck]

 

Nonmetallic chemical element, chemical symbol P, atomic number 15. The ordinary allotrope, called white phosphorus, is a poisonous, colourless, semitransparent, soft, waxy solid that glows in the dark (see phosphorescence) and combusts spontaneously in air, producing dense white fumes of the oxide P4O10; it is used as a rodenticide and a military smokescreen. Heat or sunlight converts it to the red phosphorus allotrope, a violet-red powder that does not phosphoresce or ignite spontaneously. Much less reactive and soluble than white phosphorus, it is used in manufacturing other phosphorus compounds and in semiconductors, fertilizers, safety matches, and fireworks. Black phosphorus, made by heating the white form under pressure, is flaky like graphite. Phosphorus seldom occurs uncombined in nature. As the phosphate ion, it is abundant and widely distributed, in apatite, phosphorite, and many other minerals. Phosphorus has valence 3 or 5 in compounds, which have many uses in industry. Phosphine (PH3) is a chemical raw material and a doping agent (deliberately added impurity) for solid-state electronics components. Organic phosphorus compounds are used as plasticizers, gasoline additives, insecticides (e.g., parathion), and nerve gases. In living organisms the role of phosphorus is essential; it is a component of DNA and RNA, ATP, and bone.

For more information on phosphorus, visit Britannica.com.

 

An essential macronutrient that constitutes approximately 22% of the body's total mineral content. Most of the phosphorus is combined with calcium in bones. Phosphorus is a constituent of many vital compounds including DNA, ATP, and phospholipids. Dietary deficiency is rare.

 
(fŏs'fərəs) [Gr.,=light-bearing], nonmetallic chemical element; symbol P; at. no. 15; at. wt. 30.97376; m.p. 44.1°C; b.p. about 280°C; sp. gr. 1.82 at 20°C; valence −3, +3, or +5. Solid phosphorus has a tetratomic molecule (P4) with molecular weight 123.8952 atomic mass units (amu). Phosphorus was discovered c.1674 by Hennig Brand of Hamburg, an alchemist, who prepared it from urine. Phosphoric acid was discovered in 1770 by K. W. Scheele and J. G. Gahn in bone ash (see ash); Scheele later isolated phosphorus from bone ash (1774) and produced phosphoric acid by the action of nitric acid on phosphorus (1777).

Forms

Phosphorus exhibits allotropy (i.e., it has multiple forms in the same physical state); the physical constants given above are for the common white phosphorus. White phosphorus is an extremely poisonous, yellow to white, waxy, solid substance, nearly insoluble in water but very soluble in carbon disulfide. When exposed to air it ignites spontaneously, burning to form white fumes of phosphorus pentoxide, P2O5. Because of its toxicity and pyrophoric nature, phosphorus is stored underwater. Contact with the skin may cause burns. White phosphorus is phosphorescent (i.e., glows without emitting heat).

When white phosphorus is heated to about 250°C in the absence of air, it changes into the more stable red phosphorus. This form appears as dull, reddish-brown cubic crystals or amorphous powder. Its specific gravity is 2.34. The red form is less dangerous than the white form, but should be handled with caution. It is insoluble in carbon disulfide and most other solvents. It does not ignite unless heated to about 200°C, does not phosphoresce, and is not poisonous. Another form of phosphorus is black phosphorus, a crystalline electrically conductive material similar to graphite in appearance. It was first prepared by P. W. Bridgman by heating white phosphorus to 200°C under a pressure of 12,000 atmospheres. Its specific gravity is 2.70.

Natural Occurrence and Commercial Preparation

Because of its chemical activity phosphorus does not occur uncombined in nature but is widely distributed in many minerals. A major source is apatite, an impure calcium phosphate mineral found in phosphate rocks. In the United States major deposits are found in Florida, Tennessee, Montana, and Idaho. White phosphorus is prepared commercially from phosphate rock in an electric furnace or blast furnace. The principal use of phosphorus is in compounds; for this reason, most of the phosphorus produced in furnaces is burned to make phosphorus pentoxide, a white powdery substance. While the pentoxide is used as a drying agent and chemical reagent, it is chiefly converted to phosphoric acid, H3PO4, also called orthophosphoric acid, by reaction with water. Another important source of phosphoric acid is from phosphate rocks by treatment with sulfuric acid; this is the so-called wet-acid process.

Biological Importance and Applications

Phosphorus is present in plants and animals. There is over 1 lb (454 grams) of phosphorus in the human body. It is a component of adenosine triphosphate (ATP), a fundamental energy source in living things. It is found in complex organic compounds in the blood, muscles, and nerves, and in calcium phosphate, the principal material in bones and teeth. Phosphorus compounds are essential in the diet. Organic phosphates, ferric phosphate, and tricalcium phosphate are added to foods. Dicalcium phosphate is added to animal feeds.

White phosphorus is used as a deoxidizing agent in the preparation of steel and phosphor bronze. It is also used in rat poisons and to make smoke screens (by burning) for warfare. Red phosphorus is used in making matches. The major use of phosphorus compounds is in fertilizers, especially in a mixture called superphosphate, obtained from phosphate minerals by sulfuric acid treatment; and in nitrophosphates. Phosphorus compounds are also used commercially in detergents, water softeners, pharmaceuticals, dentifrices, and in many other less important uses. Toxic nerve gases such as sarin contain phosphorus.

Phosphoric acid is primarily used in the production of phosphate compounds. It is also used in pickling metals, in sugar refining, and in soft drinks. Phosphorus forms a number of compounds with the halogens, e.g., the trichloride, PCl3, and the pentachloride, PCl5, both used as reagents. It also forms an oxychloride, POCl3. It reacts with sulfur to form a pentasulfide, P2S5, and a thiochloride, PSCl3, used in insecticides and oil additives. Phosphine, PH3, is a poisonous gas. Besides the pentoxide, phosphorus forms several other oxides; there are several acids other than the orthophosphoric acid noted above. Phosphorus also combines with various other nonmetals and with some metals.


 

A chemical element, atomic number 15, atomic weight 30.974, symbol P. Phosphorus is an essential element in the diet. In the form of phosphates it is a major component of the mineral phase of bone and is involved in almost all metabolic processes. It also plays an important role in cell metabolism. It is obtained by the body from milk products, cereals, meat and fish, and its use by the body is controlled by vitamin D and calcium.

  • p.-32 — a radioisotope of phosphorus having a half-life of 14.3 days and emitting only beta rays; used in the form of sodium phosphate P-32 for treatment of polycythemia vera, chronic myelocytic leukemia and chronic lymphocytic leukemia, and in localizing certain tumors during surgery. Symbol 32P.
  • calcium:p. ratio — see calcium: phosphorus ratio.
  • inorganic p. — any phosphorus-containing compound which does not also contain carbon.
  • p. nutritional deficiency — causes rickets in the young and osteomalacia in adult ruminants. In less severe deficiency states there is pica, growth retardation, infertility and possibly retention of placenta. See also postparturient hemoglobinuria. An unlikely nutritional deficiency in carnivores.
  • p. poisoning — is very rare because of the absence of elemental phosphorus from the environment. Causes severe gastroenteritis with vomiting and diarrhea. If the animal survives the gastroenteritis there is a subsequent acute hepatic insufficiency.
  • p. restriction — indicated in the dietary management of chronic renal disease and secondary hyperaparathyroidism; in dogs and cats, usually accomplished by reducing the content of meat.
  • p. supplements — supplementing the diets of animals exposed to phosphorus deficient feeds is usually achieved by feeding bone meal, or calcium or sodium phosphates. All are readily assimilable but none are palatable and special devices are often necessary to get animals to take required amounts. See also dietary phosphate.
 

The chemical symbol for phosphorus. Used in the formula NPK of a complete fertilizer.

 
Wikipedia: phosphorus
15 siliconphosphorussulfur
N

P

As
P-TableImage.png
General
Name, symbol, number phosphorus, P, 15
Chemical series nonmetals
Group, period, block 153, p
Appearance waxy white/ red/
black/ colorless
P,15.jpg
Standard atomic weight 30.973762(2) g·mol−1
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)
Heat of fusion (white) 0.66 kJ·mol−1
Heat of vaporization 12.4 kJ·mol−1
Heat capacity (25 °C) (white)
23.824 J·mol−1·K−1
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
Oxidation states ±3, 5, 4
(mildly acidic oxide)
Electronegativity 2.19 (Pauling scale)
Ionization energies
(more)
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
Miscellaneous
Magnetic ordering no data
Thermal conductivity (300 K) (white)
0.236 W·m−1·K−1
Bulk modulus 11 GPa
CAS registry number 7723-14-0
Selected isotopes
Main article: Isotopes of phosphorus
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 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 Star").

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 fertilisers.

Phosphorus compounds are also widely used in explosives, nerve agents, friction matches, fireworks, pesticides, toothpaste, and detergents.

Characteristics and allotropes

Elemental phosphorus can exist in several allotropes, most commonly white, red and black.

White phosphorus (P4) exists as individual molecules made up of four atoms in a tetrahedral arrangement, resulting in very high ring strain and instability. It contains 6 single bonds.

White_phosphrous_molecule.jpg

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 disulfide.

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 silica[1]. Elemental 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 urine.

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 forms.

Hittoff_phosphorus_chain.jpg

One of the forms of red/black phosphorus is a cubic solid.[2]

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.[3][4] A recent synthesis of black phosphorus using metal salts as catalysts has been reported.[5]

Black_phosphorus.jpg

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.[6] 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,[7] there is only a range of partial pressure where it does. Heat can be applied to drive the reaction at higher pressures.[8]

In 1974 the glow was explained by R. J. van Zee and A. U. Khan.[6] 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 is properly called luminescence (glowing by its own reaction, in this case chemoluminescence), not phosphorescence (re-emitting light that previously fell on it).

Applications

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 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 framework of these molecules. Living cells also use 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 macromineral, 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.

History

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[1]. The precipitated phosphates were made from ground-up bones that had been de-greased and treated with strong acids[1]. This process became obsolete in the late 1890s when the electric arc furnace was adapted to reduce phosphate rock[1].

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.)[6] 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.[6][1] In World War I it was used in incendiaries, smoke screens and tracer bullets[1]. A special incendiary bullet was developed to shoot at hydrogen-filled Zeppelins over Britain (hydrogen being highly inflammable if it can be ignited)[1]. 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,[9] 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.

Occurrence

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[1]. 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.[10]

See also Phosphate minerals.

Precautions

Hazard_F.svg
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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". [11]

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 hemolysis."[12]

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 Phosphorus injury.[13]

DEA List I status

Phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine.[14] 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.[15] 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 substances.[15][16][17]

As an exception to the octet rule

For more details on this topic, see Octet rule.

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 rule.

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.

Isotopes

For more details on this topic, see Isotopes of phosphorus.

Radioactive isotopes of phosphorus include

  • 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, Lucite, plastic, wood, or water.[18]
  • 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.

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, phosphorus forms phosphorous and phosphoric compounds.

Compounds

See also Phosphorus compounds

References

  1. ^ a b c d e f g h i j k l m n
  2. ^ R. Ahuja, Physica Status Solidi, Sectio B: Basic Research, 2003, 235, 282-287
  3. ^ A. Brown, S. Runquist, Acta Crystallogr., 19 (1965) 684
  4. ^ Cartz, L.;Srinivasa, S.R.;Riedner, R.J.;Jorgensen, J.D.;Worlton, T.G., Journal of Chemical Physics, 1979, 71, 1718-1721
  5. ^ Stefan Lange, Peer Schmidt, and Tom Nilges, Inorganic Chemistry, 2007, 46, 4028
  6. ^ a b c d Emsley, John (2000). The Shocking History of Phosphorus. London: Macmillan. ISBN 0-330-39005-8
  7. ^ Nobel Prize in Chemistry 1956 - Presentation Speech, by Professor A. Ölander (committee member)
  8. ^ Phosphorus Topics page, at Lateral Science
  9. ^ Aall C. H. (1952). "The American Phosphorus Industry". Industrial & Engineering Chemistry 44. DOI:10.1021/ie50511a018. 
  10. ^ Podger, Hugh, (2002). Albright & Wilson: The Last 50 Years. Studley: Brewin Books. ISBN 1-85858-223-7
  11. ^ emedicine.com CBRNE - Incendiary Agents, White Phosphorus (Smoking Stool Syndrome)
  12. ^ US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries
  13. ^ Emergency War Surgery NATO Handbook: Part I: Types of Wounds and Injuries: Chapter III: Burn Injury: Chemical Burns And White Phosphorus injury.
  14. ^ Skinner (1990). Methamphetamine Synthesis Via Hydriodic Acid/Red Phosphorus Reduction of Ephedrine. Forensic Sci. Int'l, 48, 123-34.
  15. ^ a b 66 FR 52670—52675. 17 October 2001.
  16. ^ 21 CFR 1309