phosphorus

 

(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.

15 silicon ← phosphorus → sulfur N ↑ P ↓ As Periodic table - Extended periodic table
General
Name, symbol, number	phosphorus, P, 15
Chemical series	nonmetals
Group, period, block	15, 3, p
Appearance	waxy white/ red/ black/ colorless
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 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.

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]

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 (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.
• 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 sodium tripolyphosphate and other sodium phosphates^[1]. Among other uses these are used to improve the characteristics of processed meat and cheese. Others are used in toothpaste^[1]. Trisodium phosphate is used in cleaning agents to soften water and for preventing pipe/boiler tube corrosion.
• Phosphorus is widely used to make organophosphorus compounds, through the intermediates phosphorus chlorides and the two phosphorus sulfides: phosphorus pentasulfide, and phosphorus sesquisulfide.^[1] Organophosphorus compounds have many applications, including in plasticizers, flame retardants, 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 dopant for 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 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

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

• Ammonium phosphate ((NH₄)₃PO₄)
• Calcium phosphate (Ca₃(PO₄)₂)
• Calcium dihydrogen phosphate (Ca(H₂PO₄)₂)
• Calcium phosphide (Ca₃P₂)
• Iron(III) phosphate (FePO₄)
• Iron(II) phosphate (Fe₃(PO₄)₂)
• Gallium(III) phosphide (GaP)
• Hypophosphorous acid (H₃PO₂)
• Lawesson's reagent
• Parathion
• Phosphine (Phosphorus Trihydride PH₃)
• Phosphoric acid (H₃PO₄)
• Phosphorus pentabromide (PBr₅)
• Phosphorus pentasulfide (P₂S₅)
• Phosphorus pentoxide (P₄O₁₀)
• Phosphorus sesquisulfide (P₄S₃)
• Phosphorus tribromide (PBr₃)
• Phosphorus trichloride (PCl₃)
• Phosphorus triiodide (PI₃)
• Sarin
• Soman
• Tabun
• Triphenyl phosphine
• Monopotassium phosphate (KH₂PO₄)
• Trisodium phosphate (Na₃PO₄)
• VX nerve gas

See also Phosphorus compounds

References

• Los Alamos National Laboratory – Phosphorus

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