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Phosphorus
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Everything about Phosphorus totally explained » see Phosphorus (morning star) for the Greek name of the morning star.
Phosphorus,, is the chemical element that has the symbol P and atomic number 15. The name comes from the (meaning "light") and φόρος (meaning "bearer"). 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 on Earth. One form of phosphorus (white phosphorus) emits a faint glow upon exposure to oxygen — hence its Greek derivation, Φωσφόρος meaning "light-bearer" (Latin Lucifer), the planet Venus as " 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 fertilizers.
Phosphorus compounds are also widely used in explosives, nerve agents, friction matches, fireworks, pesticides, toothpaste, and detergents.
Characteristics
Allotropes
Phosphorus is an excellent example of an element that exhibits allotropy, as its various allotropes have strikingly different properties.
The two most common allotropes are white phosphorus and red phosphorus. A third form, scarlet phosphorus, is obtained by allowing a solution of white phosphorus in carbon disulfide to evaporate in sunlight. A fourth allotrope, black phosphorus, is obtained by heating white phosphorus under very high pressures (12,000 atmospheres). In appearance, properties and structure it's very like graphite, being black and flaky, a conductor of electricity and has puckered sheets of linked atoms. Another allotrope is diphosphorus - which is highly reactive.
White phosphorus 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 phosphorus is a white, waxy transparent solid. This allotrope is thermodynamically unstable at normal condition and will gradually change to red phosphorus. This transformation, which is accelerated by light and heat, makes white phosphorus almost always contain some red phosphorus and appear yellow. For this reason, it's 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 infamous incindiary bomb Napalm relies, among others, on this principle to spontaneously ignite. 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 P 4O 10 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
Black phosphorus has an orthorhombic structure (C mca) and is the least reactive allotrope. It consists of many six-membered rings which are interlinked. Each atom is bonded to three other atoms. A recent synthesis of black phosphorus using metal salts as catalysts has been reported.
The diphosphorus allotrope (P 2) can be obtained normally only under extreme conditions (for example, from P 4 at 1100 kelvin). Nevertheless, some advancements were obtained in generating the diatomic molecule in homogeneous solution, under normal conditions with the use by some transitional metal complexes (based on, for example, tungsten and niobium).
Glow
The glow from phosphorus was the attraction of its discovery around 1669, but the mechanism for that glow wasn't 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's oxygen being consumed. By the 18th century, it was known that in pure oxygen phosphorus doesn't glow at all; there's only a range of partial pressure at which it does. 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.
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.
Occurrence » See also .
Due to its reactivity with air and many other oxygen-containing substances, phosphorus isn't found free in nature but it's 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. About 50 per cent of the global phosphorus reserves are in the Arab nations.(External Link ) 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
At today's rate of consumption, the supply of phosphorus is estimated to run out in 345 years.
Compounds
See also
Hydride: PH3
Halides: PBr5, PBr3, PCl3, PI3
Oxides:P4O6, P4O10
Sulfides: P2S5, P4S3
Acids: H3PO2, H3PO4
Phosphates: (NH4)3PO4, Ca3(PO4)2), FePO4, Fe3(PO4)2, Na3PO4, Ca(H2PO4)2, KH2PO4
Phosphides: Ca3P2, GaP, Zn3P2
Organophosphorus and organophosphates: Lawesson's reagent, Parathion, Sarin, Soman, Tabun, Triphenyl phosphine, VX nerve gas
As an exception to the 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 doesn't contribute to the electron density on P.
However, it should always be remembered that the octet rule is a not some universal rule of chemical bonding, and while many compounds obey it, there are many elements (the majority, in fact) to which it just doesn't apply.
Applications
Phosphorous being an essential plant nutrient find its major use as a constituent of fertilizers for agriculture and farm production in the form of concentrated phosphoric acids, which can consist of 70% to 75% P2O5. Global demand for fertilizers led to large increase in phosphate (PO43-) production in the second half of the 20th century.
Phosphorus is widely used to make organophosphorus compounds, through the intermediates phosphorus chlorides and the two phosphorus sulfides: phosphorus pentasulfide, and phosphorus sesquisulfide. These include mono-calcium phosphate which is employed in baking powder and sodium tripolyphosphate and other sodium phosphates in particular the herbicide glyphosate sold under the brand name Roundup. Production of white phosphorus takes place at large facilities and it's 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.
Spelling and etymology
According to the Oxford English Dictionary the correct spelling of the element is phosphorus. The word phosphorous is the adjectival form of the P3+ valency: so, just as sulfur forms sulfurous and sulfuric compounds, phosphorus forms phosphorous compounds (see for example phosphorous acid) and P5+ valency phosphoric compounds (see for example Phosphoric acids and phosphates).
Precautions
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".
When the white form is exposed to sunlight or when it's heated in its own vapour to 250°C, it's transmuted to the red form, which doesn't phosphoresce in air. The red allotrope doesn't spontaneously ignite in air and isn't 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's 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."
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's 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.
DEA List I status
Phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine. 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. 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.
Further Information
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