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Phosphorus

2007 Schools Wikipedia Selection. Related subjects: Chemical elements


                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
                                           Atomic mass 30.973762 (2) g/mol
                                     Electron configuration [Ne] 3s^2 3p^3
                                               Electrons per shell 2, 8, 5
                                                       Physical properties
                                                               Phase solid
                              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)
                                                Critical temperature 994 K
                                  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

   CAPTION: Vapor pressure (white)

                                          P/Pa   1  10  100 1 k 10 k 100 k
                                         at T/K 279 307 342 388 453   549

                        CAPTION: Vapor pressure (red)

                                          P/Pa   1  10  100 1 k 10 k 100 k
                                         at T/K 455 489 529 576 635   704

                                                         Atomic properties
                                                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

                CAPTION: 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 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.

   Phosphorus exists in several allotropes, most commonly white, red and
   black. White phosphorus (P[4]) contains only four atoms, resulting in
   very high ring strain and instability. White phosphorus glows in the
   dark, is highly flammable and pyrophoric (self-igniting) upon contact
   with air as well as toxic. 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 is amorphous and is the least
   reactive allotrope.

   Red phosphorus is formed by heating white phosphorus to 250°C (482°F)
   or by exposing white phosphorus to sunlight.

   Due to its high reactivity, phosphorus is never found as a free element
   in nature. It emits a faint glow upon exposure to oxygen (hence its
   Greek derivation and the Latin meaning 'morning star') and is an
   essential element for living organisms. The most important commercial
   use of phosphorus-based chemicals is the production of fertilizers.
   They are also widely used in explosives, nerve agents, friction
   matches, fireworks, pesticides, toothpaste, and detergents.

Characteristics

   Phosphorus, in its common form, is a waxy white (or yellowish) solid
   that has a characteristic, disagreeable smell similar to that of
   garlic. Pure forms of the element are colorless and transparent. This
   nonmetal is not soluble in water, but is soluble in carbon disulfide.
   The white allotrope ignites spontaneously in air; however both white
   and red phosphorus burn in air to produce phosphorus pentoxide.

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. 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 P[2]O[2] 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).

Applications

   Concentrated phosphoric acids, which can consist of 70% to 75% P[2]O[5]
   are very important to agriculture and farm production in the form of
   fertilizers. Global demand for fertilizers led to large increases in
   phosphate (PO[4]^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 elementary phosphorus is used in food
       applications such as soda beverages. The acid is also a starting
       point to make food grade phosphates. These include mono-calcium
       phosphate which is employed in baking powder and sodium
       tripolyphosphate and other sodium phosphates. Among other uses,
       these are used to improve the characteristics of processed meat and
       cheese. Others are used in toothpaste. 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. 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,, safety matches and, most notoriously, methamphetamine.
          + Phosphorus sesquisulfide is used in heads of strike-anywhere
            matches.
     * 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).
     * Red phosphorus is used in cap gun caps.

Biological role

   Phosphorus is a key element in all known forms of life. Inorganic
   phosphorus in the form of the phosphate PO[4]^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 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, see eutrophication and
   algal blooms.

History

   Phosphorus ( Greek phosphoros 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.

   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
   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 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
   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 (see precautions below). 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.

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 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 business.

   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

   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 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. 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.

   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) 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.

Isotopes

   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, acrylic, Lucite, plastic, wood, or
       water.

     * ^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 P^3+ valency: so, just as sulfur forms sulfurous and sulfuric
   compounds, so phosphorus forms phosphorous and phosphoric compounds.

Compounds

     * Ammonium phosphate ((NH[4])[3]PO[4])
     * Calcium phosphate (Ca[3](PO[4])[2])
     * Calcium dihydrogen phosphate (Ca(H[2]PO[4])[2])
     * Calcium phosphide (Ca[3]P[2])
     * Iron(III) phosphate (FePO[4])
     * Iron(II) phosphate (Fe[3](PO[4])[2])
     * Gallium(III) phosphide (GaP)
     * Hypophosphorous acid (H[3]PO[2])
     * Lawesson's reagent
     * Parathion
     * Phosphine (Phosphorus Trihydride PH[3])
     * Phosphoric acid (H[3]PO[4])
     * Phosphorus pentabromide (PBr[5])
     * Phosphorus pentasulfide (P[2]S[5])
     * Phosphorus pentoxide (P[2]O[5])
     * Phosphorus sesquisulfide (P[4]S[3])
     * Phosphorus tribromide (PBr[3])
     * Phosphorus trichloride (PCl[3])
     * Phosphorus triiodide (PI[3])
     * Sarin
     * Soman
     * Tabun
     * Triphenyl phosphine
     * Monopotassium phosphate (KH[2]PO[4])
     * Trisodium phosphate (Na[3]PO[4])
     * VX nerve gas

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