Introduction to Phosphorus
The element phosphorus, situated just below nitrogen in Group 15 of the Periodic Table, is a highly reactive non-metal. As such, it cannot exist as a free element on Earth. Phosphorus assumes various chemical states called “allotropes;” these include white, red, violet, and black phosphorus. The most widely encountered forms include white and red. It is essential to life, and has bountiful use as a fertilizer and in matches.
Cool Facts About Phosphorus
- The name “phosphorus” derived from the Greek term “phosphoros,” which translates to “bringer of light”
- Texts often refer to phosphorus as the “Devil’s Element” because of its glow, tendency to combust, and status as the 13th discovered element
- The average adult body contains about 750 grams of phosphorus–this element facilitates all forms of life
- Our bones and teeth store about 85% of the phosphorus in our bodies
- Militaries use white phosphorus to make incendiary devices such as bombs, grenades and rockets
- Phosphorus comprises around 12% of all commercial fertilizer
- Over half of phosphorus consumption occurs in the developing world
- Foods such as dairy, eggs, meat, poultry, fish, nuts, and legumes contain the highest amounts of phosphorus
Phosphorus on the Periodic Table
Phosphorus, a highly reactive waxy nonmetal, has symbol P and atomic number 15. It lies to the left of sulfur, to the right of silicon, and below nitrogen in group 15 of the periodic table, the pnictogens. The phosphorus atom has an electron configuration of [Ne]3s23p3 or 1s22s22p63s23p3
Allotropes of Phosphorus
White phosphorus is the most toxic allotrope of this element. It appears as a waxy, translucent solid substance that fluoresces in the dark. When exposed to air, white phosphorus becomes spontaneously flammable, and contact with skin causes severe burns. When white phosphorus burns, it forms phosphorus pentoxide.
In white phosphorus molecules, four phosphorus atoms connect to one another in a closed ring structure via covalent bonds. The low bond angle that results from this configuration generates strain within the molecule, explaining its extremely reactive qualities.
White phosphorus is extremely dangerous, and can cause great harm or death. It slowly changes to red phosphorus over time when exposed to heat or light, and often looks yellow because of this.
This allotrope is less toxic than its counterpart. It appears as a red-iron, lustrous, powdered substance. Red phosphorus is much more stable and less dangerous than white phosphorus. It can be extracted in small quantities from striker strips found on match boxes.
Red phosphorus forms when white phosphorus rings polymerize via covalent bonds to form straight chains. These chains minimize intramolecular strain, reducing the reactivity of the molecule.
Black phosphorus is even less reactive than red phosphorus, and it appears as a black, lustrous, crystalline solid. These molecules contain crystal lattices that form through the linkage of existing P-P-P bonds.
The larger bond angle resulting from this configuration contributes to its relative stability. As a less reactive allotrope, black phosphorus does not display any reactivity toward sulfur, oxygen, or halogens.
Black phosphorus can assume either an α-black or β-black form. The former, more stable, occurs through the heating of red phosphorus. The latter, more reactive, synthesizes after heating white phosphorus under pressure.
Violet phosphorus, the least reactive allotrope, reacts slowly with halogens. It appears as a nearly all-black, but slightly purple crystalline solid.
Violet phosphorus forms via heating red phosphorus in a sealed tube or dissolving white phosphorus in molten lead. Scientists continue to investigate its lattice structure using x-ray diffraction methods.
Phosphorus in the Environment
Because phosphorus cannot naturally occur in the environment, it is typically mined from phosphates. The term “phosphates” describes any chemical derivative of phosphorus that contains the phosphate ion, PO43-.
Phosphates typically accumulate at the bottom of rivers and lakes, within different rocks, and in large mineral deposits around the world. Natural phosphates exist inside every living thing. They comprise DNA, RNA, ATP, phospholipids, and other compounds crucial for cellular function and, by extension, life.
Once synthesized from phosphates, phosphorus holds many uses. Phosphorus facilitates the production of phosphor bronze and steel, crucial components of springs, bolts, and electrical switches. Many laundry and dishwasher detergents include phosphorus because of its ability to strip grime from objects. LED lights rely on the fluorescent property of white phosphorus. Red phosphorus is used in matches, incendiary shells, and other explosive devices due to its combustible properties.
History of the Element Phosphorus
German physician and alchemist Hennig Brand discovered the element phosphorus in 1669 by evaporating urine, then heating the residue to create phosphorus vapor. He then distilled this vapor by condensing it in water, creating solid phosphorus in the white allotropic form.
Brandt stumbled upon phosphorus in his attempts to synthesize the “Philosopher’s Stone,” which could allegedly turn metals into gold. His pursuit of this goal had involved heating sand and charcoal, then combining them with boiled urine. On one particular day, the resulting mixture produced this white vapor. He named the glowing, waxy drops “phosphorus,” a Latin term for substances that emit light.
Phosphorus – Reactions, Compounds, Oxidation States, and Isolation
Phosphorus + Oxygen
This element reacts readily with oxygen to combust, forming dense white fumes.
P4(s) + 5O2(g) → P4O10(g)
Phosphorus + Halogens
Phosphorus reacts with all halogens to form phosphorus trihalides. Specifically, it reacts with fluorine: F2, chlorine: Cl2, bromine: Br2, and iodine: I2, to form respectively phosphorus(III) fluoride: PF3, phosphorus(III) chloride: PCl3, phosphorus(III) bromide: PBr3, and phosphorus(III) iodide: PI3.
Fluorine: P4(s) + 6F2(g) → 4PF3(g)
Chlorine: P4(s) + 6Cl2(g) → 4PCl3(l)
Bromine: P4(s) + 6Br2(g) → 4PBr3(l)
Iodine: P4(s) + 6I2(g) → 4PI3(g)
Phosphorus + Water
Because phosphorus does not react with water, scientists often store it underwater to prevent reactions with air. But boiling phosphorus in water produces phosphine and phosphorous acid.
2P4(s) + 3H2O(l) → PH3(g) + H3PO3(s)
Phosphorus + Acids
Phosphorus reacts with HCl to produce phosphine and phosphorus(III) chloride.
P4(s) + 6HCl(l)→ 2PH3(g) + 2PCl3(l)
Phosphorus exists used almost entirely in the form of compounds, which typically utilize its +5 oxidation state.
Common phosphate compounds include calcium phosphate: Ca3(PO4)2, apatite: Ca5(PO4)3OH, fluorapatite: Ca5(PO4)3F; and chlorapatite: Ca5(PO4)3Cl. The most popular phosphate fertilizer contains the compound calcium dihydrogen phosphate, or superphosphate: Ca(H2PO4)2.
The catalyst of many organic phosphorus synthesis reactions is the highly toxic compound hydrogen phosphine: PH3. The reaction of white phosphorus with sodium or potassium hydroxide produces phosphine.
Phosphorus pentoxide: P4O10, which forms when white phosphorus is burned, functions as a dehydrating and condensing agent. It reacts vigorously with water to form phosphoric acid. Phosphorus is in the +5 oxidation state.
Tetraphosphorus hexaoxide: P4O6 functions as the anhydride of phosphorous acid. Phosphorus is in the +3 oxidation state in this compound.
Treating phosphorus pentoxide with water generates orthophosphoric acid: H3PO4, which has industrial uses. Phosphorous acid: H₃PO₃ provides the intermediate during the preparation of other phosphorus compounds.
Phosphorus Oxidation States
The most common oxidation number of phosphorus is +5. Oxidation states of +3 and -3 also occur, in phosphites and phosphides respectively. Phosphides form from the reaction of metals with red phosphorus. Phosphine also contains phosphorus in the +3 state.
Isolation of Phosphorus
Scientists typically isolate elemental phosphorus via the following reaction, using calcium phosphate obtained from phosphate rock. This reaction occurs at 1500 degrees Celsius. Aluminum can also be used, which allows the reaction to take place at a lower temperature.
2Ca3(PO4)2(s) + 6SiO2(s) + 10C(s) → 6CaSiO3(l) + 10CO(g) + P4(g)
Phosphorus’ Applications to Today’s World
What are the uses of Phosphorus?
The main use of phosphorus, in the form of concentrated phosphoric acid, in is fertilizers. Phosphorus compounds are also used in baking powder, pesticides, detergents, and plasticisers.
Phosphate is an important ingredient in fertilizers, and helps sustain high crop yields. The entire agriculture industry depends on it, and it has no easy substitute. However, phosphate rock is being mined at unsustainable levels. Some researchers are trying to recover phosphorus from wastewaters, by using bacteria to precipitate the phosphorus.
Phosphorus can also facilitate the synthesis of d-methamphetamine, one of the two stereoisomers of the compound methamphetamine. Red phosphorus can react with iodine in the presence of water to yield precursors needed for the drug. For this reason, the possession and distribution of both red and white phosphorus (which can be converted to the red allotrope of the element) with the intent of generating methamphetamine remains illegal in the United States.
The element phosphorus also contributes to eutrophication, or decreased oxygen levels in aquatic habitats. Lakes, streams, and rivers near facilities that process phosphorus often receive deposits of this substance. This prompts the overgrowth of algae that depletes oxygen from the water. Water runoff carrying phosphorus-containing fertilizer also pool to create these “dead zones.” Scientists combat eutrophication by limiting the amount of phosphorus contained in fertilizers and detergents, and by maintaining clean aquatic environments.
Phosphorus Physical Properties
- Atomic Symbol: P
- Melting point: 317.25 K; 44.10°C; 111.40°F
- Boiling point: 553.15 K; 280.00 °C; 536.00 °F
- Density: 1.82 g/cm^3
- Atomic weight: 30.97
- Atomic number: 15
- Electronegativity: 2.19
- Classification: non-metal
- Natural abundance of X% in the earth’s crust: 0.07%
- Electron shell configuration: [Ne] 3s²3p³
- Isotopes: Phosphorus has 21 recognized isotopes with mass numbers spanning from 24-46
- Only P-31 is stable and naturally occurring
- P-26 through P-30; P-30 through P-46 are artificially produced and radioactive
- Found: Because of its highly reactive behavior, phosphorus does not occur in its pure elemental form on Earth; it is primarily found in minerals called phosphates, such as apatite and fluorapatite
- Toxicity: All forms of the element phosphorus are extremely toxic to humans, but the most toxic form is white phosphorus
Where Can I Buy Phosphorus?
Violet phosphorus, and sometimes white phosphorus, is available for purchase online in very small amounts.
Elemental phosphorus is refined and isolated from phosphate compounds found in bone ash, natural minerals, or even urine. The refinement process is detailed here; this should only be performed in a lab with professional aid and supervision.
In many countries, including the USA, the sale of red phosphorus is heavily restricted, because it will reduce elemental iodine to hydroiodic acid. HI can then be used to reduce ephedrine or pseudoephedrine to methamphetamine. Sigma Aldrich sells red phosphorus to its screened list of customers. People have also been known to purchase it from overseas suppliers.