Phosphorus

Characteristics

Allotropes

Main article: Allotropes of phosphorus

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 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 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, 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.

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

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

The diphosphorus allotrope (P₂) can be obtained normally only under extreme conditions (for example, from P₄ at 1100 kelvin). Nevertheless, some advancements were obtained in generating the diatomic molecule in homogenous solution, under normal condtitions with the use by some transitional metal complexes (based on for example niobium).^[6]

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.^[7] It was known from early times that the glow would persist for a time in a stoppered jar but then cease. partial pressure where it does. Heat can be applied to drive the reaction at higher pressures.^[9]

In 1974 the glow was explained by R. J. van Zee and A. U. Khan.^[7] 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 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 (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, 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 elemental phosphorus is used in food applications such as soda beverages. The acid is also a starting point to make food grade phosphates.^[1] These include mono-calcium phosphate which is employed in baking powder and corrosion.
• Phosphorus is widely used to make 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,^[1] safety matches, pharmaceutical grade and street methamphetamine, and is used in cap gun caps.
• Phosphorus sesquisulfide is used in heads of strike-anywhere matches.^[1]
• In trace amounts, phosphorus is used as a N-type semiconductors.
• ³²P and ³³P are used as radioactive tracers in biochemical laboratories (see Isotopes).

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 use phosphate to transport cellular energy via Calcium phosphate salts assist in stiffening bones.

An average adult human contains a little less than 1 kg of phosphorus, about 85% of which is present in bones and teeth in the form of apatite, and the remainder inside cells in soft tissues. A well-fed adult in the industrialized world consumes and excretes about 1-3 g of phosphorus per day in the form of phosphate. Only about 0.1% of body phosphate circulates in the blood, but this amount reflects the amount of phosphate available to soft tissue cells.

In medicine, low phosphate syndromes are caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes which draw phosphate from the blood or pass too much of it into the urine. All are characterized by hypophosphatemia (see article for medical details). Symptoms of low phosphate include muscle and neurological dysfunction, and disruption of muscle and blood cells due to lack of ATP.

Phosphorus is an essential 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 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 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).^[7] 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.^[7][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 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,^[10] 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

See also Phosphate minerals.

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.^[11]

At today's rate of consumption, the supply of phosphorous is estimated to run out in 345 years.^[12]

Precautions

Organic compounds of phosphorus form a wide class of materials, some of which are extremely toxic. Fluorophosphate 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". ^[13]

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."^[14]

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.^[15]

DEA List I status

Phosphorus can reduce elemental List I precursor chemicals under 21 CFR 1310.02 effective November 17, 2001.^[17] 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.^[17][18][19]

As an exception to the octet rule

For more details on this topic, see Octet rule.

The simple octet rule.

An alternate description of the bonding, however, respects the octet rule by using electron density on P.

Isotopes

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

isotopes of phosphorus include

• ³²P; a plastic, wood, or water.^[20]
• 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} Threlfall, R.E., (1951). 100 years of Phosphorus Making: 1851 - 1951. Oldbury: Albright and Wilson Ltd
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. ^ [1]
7. ^ ^{a b c d} Emsley, John (2000). The Shocking History of Phosphorus. London: Macmillan. ISBN 0-330-39005-8
8. ^ Nobel Prize in Chemistry 1956 - Presentation Speech, by Professor A. Ölander (committee member)
9. ^ Phosphorus Topics page, at Lateral Science
10. ^ Aall C. H. (1952). "The American Phosphorus Industry". Industrial & Engineering Chemistry 44. doi:10.1021/ie50511a018.
11. ^ Podger, Hugh, (2002). Albright & Wilson: The Last 50 Years. Studley: Brewin Books. ISBN 1-85858-223-7
12. ^ (May 26, 2007) "How Long Will it Last?". New Scientist 194 (2605): 38-39. ISSN 4079 0262 4079.
13. ^ emedicine.com CBRNE - Incendiary Agents, White Phosphorus (Smoking Stool Syndrome)
14. ^ US Navy's Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries
15. ^ Emergency War Surgery NATO Handbook: Part I: Types of Wounds and Injuries: Chapter III: Burn Injury: Chemical Burns And White Phosphorus injury.
16. ^ Skinner (1990). Methamphetamine Synthesis Via Hydriodic Acid/Red Phosphorus Reduction of Ephedrine. Forensic Sci. Int'l, 48, 123-34.
17. ^ ^{a b} 66 FR 52670—52675. 17 October 2001.
18. ^ 21 CFR 1309
19. ^ 21 USC, Chapter 13 (Controlled Substances Act)
20. ^ http://www.oseh.umich.edu/TrainP32.pdf