Author Topic: Phosphorus  (Read 2766 times)

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  • Guest
« on: December 03, 2003, 04:31:00 AM »

Originally, phosphorus was extracted from urine. However there is plenty of phosphorus in phosphate ores and those ores represent the usual source for commercially produced phosphorus. There is normally no need to make phosphorus in the laboratory as it is readily available commercially.

The usial route involves heating a phosphate with sand and carbon in an electric furnace. It is highly energy intensive.

2Ca3(PO4)2 + 6SiO2 + 10C (1500°C)  6CaSiO3 + 10CO + P4

The reaction may proceed via "phosphorus pentoxide", P4O10.

2Ca3(PO4)2 + 6SiO2 +  6CaSiO3 + P4O10

P4O10 + 10C  10CO + P4

  Very good info on physical properties.

The term Red-phosphorus (P-red) is used for describing one of the allotropic forms of Phosphorus, obtained by heating White Phosphorus (P-w) at a temperature close to 300C in absence of oxygen.

The colour ranges from the orange to the dark violet depending on molecular weight, particle size and impurities. P-red is largely amorphous inorganic polymer, although X-rays have established the existence of several crystalline forms, normally present in a limited extend (<10%w).

Phosphorus exists in four or more allotropic forms: white (or yellow), red, and black (or violet). Ordinary phosphorus is a waxy white solid; when pure it is colorless and transparent. White phosphorus has two modifications: alpha and beta with a transition temperature at -3.8oC.

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 17C, 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.

Found most often in phosphate rock. Pure phosphorus is obtained by heating a mixture of phosphate rock, coke, and silica to about 1450 °C.

There are several forms of phosphorus. White phosphorus is manufactured industrially, glows in the dark, is spontaneously flammable when exposed to the air above 30°C and is a deadly poison. Red phosphorus, made by gently heating white phosphorus in the absence of air to about 250°C, does not glow, is stable and is not poisonous. This is the material, mixed with powdered glass, stuck on the side of boxes of safety matches on which the matches must be struck to light them.

Phosphorus has three main allotropes: white, red and black. White phosphorus is poisonous and can spontaneously ignite when it comes in contact with air. For this reason, white phosphorus must be stored under water and is usually used to produce phosphorus compounds. Red phosphorus is formed by heating white phosphorus to 250°C (482°F) or by exposing white phosphorus to sunlight. Red phosphorus is not poisonous and is not as dangerous as white phosphorus, although frictional heating is enough to change it back to white phosphorus. Red phosphorus is used in safety matches, fireworks, smoke bombs and pesticides. Black phosphorus is also formed by heating white phosphorus, but a mercury catalyst and a seed crystal of black phosphorus are required. Black phosphorus is the least reactive form of phosphorus and has no significant commercial uses.


  • Guest
white P glows in the dark ??
« Reply #1 on: December 03, 2003, 03:29:00 PM »
White phosphorus is manufactured industrially, glows in the dark


what? how? why? really?

without radioactive decay, how does a pure element glow in the dark?  Is it "decaying" to RP and emitting a photon
in the process?? 

Origin : The name is derived from the Greek ‘phosphoros’, meaning bringer of light,
because it glows in the dark

following up:

White phosphorus is a dangerous explosion hazard when it forms a chemical reaction with many chemicals, including alkaline hydroxides, beryllium, bromine, halogens, chlorine dioxide, chlorine trifluoride, chlorosulfonic acid, copper, iron, manganese compounds, nickel, nitrates, nitrogen dioxide, oxygen, performic acid, sulfuric acid, peroxyformic acid,
chlorosulfuric acid, hologen azides,
and hexalithium disilicide.

When exposed to air
emits a green light
and gives off white fumes.


  • Guest
Re: White phosphorus is manufactured ...
« Reply #2 on: December 03, 2003, 03:56:00 PM »

White phosphorus is manufactured industrially, glows in the dark
what? how? why? really?
without radioactive decay, how does a pure element glow in the dark?  Is it "decaying" to RP and emitting a photon
in the process?? 

The glow of phosphorus was minutely studied by Boyle, who found that solutions in some essential oils (oil of cloves) showed the same character, whilst in others (oils of mace and aniseed) there was no phosphorescence. He also noticed a strong garlic-like odour, which we now know to be due to ozone. Frederick Slare noticed that the luminosity increased when the air was rarefied, an observation confirmed by Hawksbee and Homberg, and which was possibly the basis of Berzelius’s theory that the luminosity depended on the volatility of the element and not on the presence of oxygen. Lampadius, however, showed that there was no phosphorescence in a Torricellian vacuum; and other experimenters proved that oxygen was essential to the process. It depends on the partial pressure of the oxygen and also on temperature. In compressed air at ordinary temperature there is no glowing, but it may be brought about by heating. Again, in oxygen under ordinary conditions there is no phosphorescence, but if the gas be heated to 25° glowing occurs, as is also the case if the pressure be diminished or the gas diluted. It is also remarkable that many gases and vapours, e.g. Cl, Br, I, NH3, NO2, H2S, SO2, CS2 inhibit the phosphorescence.

White phosphorus is stored under water to keep it away from oxygen. I take a small piece of white phosphorus and expose it to air. I darken the room and I can see the phosphorus glow. The glow comes from the reaction of some of the phosphorus with some oxygen. The reaction is spontaneous, but it is slow. The glow comes from the increased vibration of the parts of the molecules. The oxygen molecule accelerates toward the phosphorus atoms. When the oxygen atom gets close to the electrons of the phosphorus atoms, all points of zero force shift. This causes all of the atoms to vibrate. Some of the vibrating parts oscillate with frequencies in the range of visible light.
The oscillations of electrons in molecules are more complicated than the oscillations in free atoms, because each electron is influenced by the fields of other electrons in the same atom. In any case the principle still holds, as it does in the hydrogen atom, that a photon is emitted when the frequency of the oscillation of the electron is the same as the resonant frequency of the oscillation of the photon.
In an oxygen molecule, the point of zero force for one of the electrons is determined by all of the forces in the molecule. The points of zero force and the strength of the forces keep shifting and altering, but there is just enough stability for the electron to emit an occasional photon.

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