Phosphorus

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Phosphorus, (from the Greek language Phosphoros meaning "light bearing"), 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.

Contents

1 Notable characteristics 1.1 Forms 2 Applications 3 Biological role 4 History 5 Occurrence 6 Precautions 7 Isotopes 8 References

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.

Forms

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 phosphorus 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.

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.

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 (PO₄^3-) 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 applications as incendiary bombs, smoke pots, smoke bombs and tracer bullets.
• Red phosphorus is essential for manufacturing matchbook strikers, flares, and, most notoriously, methamphetamine.
• Miscellaneous uses; used in the making of safety matches, pyrotechnics, pesticides, toothpaste, detergents, etc.

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 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.

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. There are however concerns over how long these phosphorus deposits will last. USA will deplete their deposits around 2035. China and Morocco have the largest known deposits today, but they too will eventually be depleted. During that depletion there could be a serious problem for the worlds food production since phosphorus is such an essential ingredient in fertilizers.

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 (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 and therefore 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". 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.

Isotopes

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.

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References

• Los Alamos National Laboratory – Phosphorus (http://periodic.lanl.gov/elements/15.html)