|Boiling Point: 553°K, 280°C, 536°F
Melting Point: 317.45°K, 44.3°C, 111.7°F
Electrons Energy Level: 2, 8, 5
Isotopes: 22 + 1 Stable
Heat of Vaporization: 12.129 kJ/mol
Heat of Fusion: 0.657 kJ/mol
Density: 1.82 g/cc @ 300°K
Specific Heat: 0.77 J/g°K
Atomic Radius: 1.23Å
Ionic Radius: 0.38Å
Electronegativity: 2.19 (Pauling); 2.06 (Allrod Rochow)
Vapor Pressure: 20.8 Pa @ 44.3°C
(Greek: phôs meaning "light", and phoros meaning
"bearer"; phosphoros was the ancient name for the planet Venus, but in
Greek Mythology, Hesperus and Eosphorus could be confused with Phosphorus).
Discovered by German alchemist Hennig Brand in 1669 through a preparation from urine, the
original extraction was made from about 60 pails of urine, which naturally contains
considerable quantities of disolved 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
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 electric 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 poisoning resulted from its use.
Alchemical Symbol, Phosphorus
(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.
Additional Representations of Alchemical Symbols for Phosphorus
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. 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.
1s2 2s2p6 3s2p3
Phosphorus occurs in at least 10 allotropic forms, the most common (and reactive) of which is so-called White (or Yellow) Phosphorus which looks like a waxy solid or plastic. It is very reactive and will spontaneously inflame in air so it is stored under water. The other common form of Phosphorus is Red Phosphorus which is much less reactive and is one of the components on the striking surface of a match book. Red Phosphorus can be converted to White Phosphorus by careful heating.
Commercially, Phosphorus compounds are used in the manufacture of Phosphoric Acid (H3PO4) (found in soft drinks and used in fertilizer compounding).
Due to its high reactivity, Phosphorus is never found as a free element in nature. White Phosphorus emits a faint glow upon exposure to Oxygen (hence its Greek derivation and the Latin meaning "morning star"). Elemental Phosphorus as prepared artificially exists in several allotropes, most commonly white, red and black.
White Phosphorus (P4) contains only four atoms in a tetrahedron arrangement, resulting in very high ring strain and instability. White phosphorus, is a yellow, waxy transparent solid. It glows greenishly 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 odor of combustion of this form has a characteristic garlic smell, and samples are commonly coated with white "(Di)phosphorus pentoxide", which actually consists of P4O10 tetrahedra with oxygen inserted between the Phosphorus atoms and at their verticies. White Phosphorus is insoluble in water but soluble in Carbon Disulfide, CS2.
The white allotrope can be produced using several different methods. In one process, Tricalcium 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, H3PO4. This process is similar to the first sythesis 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. 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 has an orthorhombic structure (Cmca) and is the least reactive allotrope.
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).
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 Tricalcium 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, H3PO4, business.
Concentrated Phosphoric Acids, which can consist of 70% to 75% P2O5 are very important to agriculture 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.
|Ammonium Phosphate, (NH4)3PO4||Calcium Phosphate, Ca3(PO4)2|
|Calcium Dihydrogen Phosphate, Ca(H2PO4)2||Calcium Phosphide, Ca3P2|
|Iron (II) Phosphate, Fe3(PO4)2||Gallium (III) Phosphide, GaP|
|Monopotassium Phosphate, KH2PO4||Trisodium Phosphate, Na3PO4|
|Hypophosphorous Acid, H3PO2||Phosphorous acid (H3PO3)|
|Triphenyl Phosphine, P(C6H5)3||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|
|Lawesson's Reagent, 2,4-bis(4-methoxyphenyl)-1,3,2,4-dithiadiphosphetane-2,4-disulfide|
|VX Nerve Gas, O-ethyl S-(2-diisopropylaminoethyl) methylphosphonothiolate|
|Tabun, Ethyl N,N-dimethylphosphoramidocyanidate|
|Parathion||Triphenyl Phosphine||Lawesson's Reagent|
Radioactive isotopes of phosphorus include:
|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.|
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.
|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.|
Upon exposure to elemental phosphorus, in the past it was suggested to wash the affected area with 2% Copper Sulfate, CuSO4, 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 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.
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.
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.
Atomic Radius (Å): 1.23Å
Electrochemical Equivalents: 0.23113g/amp-hr
Atomic Mass Average: 30.97376
(Gr. phosphoros, light bearing; ancient name for the planet Venus when appearing before sunrise) Discovered in 1669 by Brand, who prepared it from urine. Phosphorus exists in four or more allotropic forms: white (or yellow), red, and black (or violet). White phosphorus has two modifications: alpha and beta with a transition temperature at -3.8oC. Never found free in nature, it is widely distributed in combination with minerals. Phosphate rock, which contains the mineral apatite, an impure tri-calcium phosphate, is an important source of the element. Large deposits are found in Russia, in Morocco, and in Florida, Tennessee, Utah, Idaho, and elsewhere. Phosphorus is an essential ingredient of all cell protoplasm, nervous tissue, and bones. Ordinary phosphorus is a waxy white solid; when pure it is colorless and transparent. It is insoluble in water, but soluble in carbon disulfide. It takes fire spontaneously in air, burning to the pentoxide. It is very poisonous, 50 mg constituting an approximate fatal dose. Exposure to white phosphorus should not exceed 0.1 mg/m3 (8-hour time-weighted average - 40-hour work week). White phosphorus should be kept under water, as it is dangerously reactive in air, and it should be handled with forceps, as contact with the skin may cause severe burns. When exposed to sunlight or when heated in its own vapor to 250oC, it is converted to the red variety, which does not phosphoresce in air as does the white variety. This form does not ignite spontaneously and is not as dangerous as white phosphorus. It should, however, be handled with care as it does convert to the white form at some temperatures and it emits highly toxic fumes of the oxides of phosphorus when heated. The red modification is fairly stable, sublimes with a vapor pressure of 1 atm at 17oC, and is used in the manufacture of safety matches, pyrotechnics, pesticides, incendiary shells, smoke bombs, tracer bullets, etc. White phosphorus may be made by several methods. By one process, tri-calcium phosphate, the essential ingredient of phosphate rock, is heated in the presence of carbon and silica in an electric furnace or fuel-fired furnace. Elementary phosphorus is liberated as vapor and may be collected under phosphoric acid, an important compound in making super-phosphate fertilizers. In recent years, concentrated phosphoric acids, which may contain as much as 70 to 75% P2O5 content, have become of great importance to agriculture and farm production. World-wide demand for fertilizers has caused record phosphate production. Phosphates are used in the production of special glasses, such as those used for sodium lamps. Bone-ash, calcium phosphate, is also used to produce fine chinaware and to produce mono-calcium phosphate used in baking powder. Phosphorus is also important in the production of steels, phosphor bronze, and many other products. Trisodium phosphate is important as a cleaning agent, as a water softener, and for preventing boiler scale and corrosion of pipes and boiler tubes. Organic compounds of phosphorus are important.
Source: CRC Handbook of Chemistry and Physics, 1913-1995. David R. Lide, Editor in Chief. Author: C.R. Hammond