Phosphorus has one of the most fascinating discovery stories in chemistry. In 1669, the German alchemist Hennig Brand (c. 1630–c. 1710) sought to transform base metals into gold. By heating and distilling large quantities of human urine, he obtained a white substance that glowed in the dark and spontaneously ignited upon contact with air. He named this discovery phosphorus (from the Greek phosphoros = light-bearer). This was the first chemical element whose discoverer and date of discovery are precisely known. In 1769, Carl Wilhelm Scheele (1742–1786) and Johan Gottlieb Gahn (1745–1818) discovered that phosphorus could be extracted from bones. In 1777, Antoine Lavoisier (1743–1794) established that phosphorus is a chemical element, not a compound.
Phosphorus (symbol P, atomic number 15) is a non-metal in group 15 (formerly group VA) of the periodic table. Its atom has 15 protons, 15 electrons, and usually 16 neutrons in its only stable isotope (\(\,^{31}\mathrm{P}\)).
Phosphorus exists in several allotropic forms with very different properties. White phosphorus (P₄) is a waxy, white-yellowish solid, highly toxic, and pyrophoric (spontaneously ignites in air at around 30 °C). Density ≈ 1.82 g/cm³, melting point: 317.3 K (44.15 °C). Red phosphorus is an amorphous polymeric form, stable, non-toxic, and non-flammable at room temperature. Density ≈ 2.16 g/cm³. Black phosphorus is the thermodynamically stable form, with a layered structure similar to graphite. Density ≈ 2.69 g/cm³. Violet phosphorus (or Hittorf's phosphorus) is another, less common allotropic form.
| Isotope / Notation | Protons (Z) | Neutrons (N) | Atomic mass (u) | Natural abundance | Half-life / Stability | Decay / Remarks |
|---|---|---|---|---|---|---|
| Phosphorus-31 — \(\,^{31}\mathrm{P}\,\) | 15 | 16 | 30.973762 u | 100 % | Stable | Only stable isotope of phosphorus; essential for all terrestrial life. |
| Phosphorus-32 — \(\,^{32}\mathrm{P}\) | 15 | 17 | 31.973907 u | Not natural | 14.268 days | Radioactive β\(^-\) decaying into sulfur-32. Widely used in molecular biology as a radioactive tracer. |
| Phosphorus-33 — \(\,^{33}\mathrm{P}\) | 15 | 18 | 32.971725 u | Not natural | 25.34 days | Radioactive β\(^-\) decaying into sulfur-33. Used in biomedical research. |
| Phosphorus-30 — \(\,^{30}\mathrm{P}\) | 15 | 15 | 29.978314 u | Not natural | 2.498 minutes | Radioactive β\(^+\) and electron capture decaying into silicon-30. |
| Other isotopes — \(\,^{24}\mathrm{P}\) to \(\,^{46}\mathrm{P}\) | 15 | 9 — 31 | — (variable) | Not natural | Milliseconds to minutes | Very unstable isotopes produced artificially; nuclear physics research. |
N.B. :
Electron shells: How electrons organize around the nucleus.
Phosphorus has 15 electrons distributed across three electron shells. Its full electronic configuration is: 1s² 2s² 2p⁶ 3s² 3p³, or simplified as: [Ne] 3s² 3p³. This configuration can also be written as: K(2) L(8) M(5).
K Shell (n=1): contains 2 electrons in the 1s subshell. This inner shell is complete and highly stable.
L Shell (n=2): contains 8 electrons distributed as 2s² 2p⁶. This shell is also complete, forming a noble gas configuration (neon).
M Shell (n=3): contains 5 electrons distributed as 3s² 3p³. The 3s orbitals are complete, while the 3p orbitals contain only 3 out of 6 possible electrons, with one electron in each of the three 3p orbitals according to Hund's rule. Thus, 3 electrons are missing to saturate this outer shell.
The 5 electrons in the outer shell (3s² 3p³) are the valence electrons of phosphorus. This configuration explains its chemical properties:
By gaining 3 electrons, phosphorus forms the P³⁻ ion (oxidation state -3), a state present in metallic phosphides, thus adopting the configuration of argon [Ar].
By losing or sharing electrons, phosphorus can exhibit positive oxidation states: +3 and +5, with the latter being the most common, particularly in phosphoric acid H₃PO₄ and phosphates.
The oxidation state 0 corresponds to elemental phosphorus, which exists in several allotropic forms: white phosphorus (P₄, highly reactive and toxic) and red phosphorus (a more stable polymer).
The electronic configuration of phosphorus, with 5 electrons in its valence shell, classifies it among the pnictogens (elements of group 15). This structure gives it characteristic properties: the ability to form three covalent bonds by sharing its three unpaired 3p electrons, the possibility of expanding its valence shell to form up to five bonds using vacant 3d orbitals, and the ability to form both single and multiple bonds. Phosphorus can accept 3 electrons to achieve the stability of a noble gas, but this P³⁻ ionic state is rare due to its large size. More commonly, phosphorus shares its electrons in covalent bonds, forming essential compounds such as phosphates. Its biological importance is crucial: phosphorus is a constituent element of DNA, RNA, and ATP (the energy molecule of cells). In industrial chemistry, it is essential for the production of phosphate fertilizers, detergents, and is part of the composition of many organophosphorus compounds.
Phosphorus is a very reactive element, especially in its white form. It readily combines with oxygen (forming P₄O₁₀ and P₄O₆), halogens, and sulfur. White phosphorus must be stored under water to prevent spontaneous ignition. Phosphorus forms compounds in oxidation states -III, +III, and +V. The most important compounds include phosphates (PO₄³⁻), phosphoric acid (H₃PO₄), phosphines (PH₃), phosphorus pentoxide (P₂O₅), and organophosphorus compounds. Phosphorus can form P-O, P-N, P-C, and P-P bonds, giving rise to an extremely rich and varied chemistry.
Phosphorus is one of the six fundamental chemical elements of life (C, H, N, O, P, S). It is indispensable for all living organisms without exception. Phosphorus is a structural component of DNA and RNA (phosphate-sugar backbone), cell membranes (phospholipids), and ATP (adenosine triphosphate), the universal molecule for energy transfer in cells. It is also the main element in bones and teeth in the form of hydroxyapatite (Ca₁₀(PO₄)₆(OH)₂). Phosphorus plays a crucial role in regulating blood pH, enzyme activation, and cell signaling. In plants, phosphorus is essential for photosynthesis, root growth, and seed formation. The phosphorus cycle in ecosystems is fundamental but slow, often making this element a limiting factor for biological growth.
Phosphorus is mainly extracted from calcium phosphate deposits (phosphorite and apatite). Unlike nitrogen, which can be captured from the atmosphere, phosphorus must be mined. Global phosphate reserves are concentrated in a few countries: Morocco (over 70% of global reserves), China, Algeria, Syria, and South Africa. This geographical concentration raises issues of global food security, as phosphorus is irreplaceable for agriculture. Phosphorus has no known substitute in agriculture, and its recycling from wastewater and organic waste is becoming a major environmental challenge. The gradual depletion of high-quality phosphate deposits is a concern for future food production.
Phosphorus is produced in massive stars during the final stages of nuclear fusion, primarily through neutron capture. Supernovae disperse phosphorus into the interstellar medium. However, phosphorus is relatively rare in the universe compared to other biogenic elements such as carbon, nitrogen, or oxygen. Its cosmic rarity could be a limiting factor for the emergence of life as we know it elsewhere in the universe. Astronomers have detected phosphorus in comets, suggesting that these celestial bodies may have brought this essential element to the early Earth. The search for phosphorus compounds in exoplanets and their atmospheres could serve as an indirect biosignature.
N.B.:
White phosphorus is one of the most dangerous substances handled in chemistry. It spontaneously ignites upon contact with air at around 30 °C, producing a ghostly greenish light and toxic fumes of phosphorus pentoxide. Burns from white phosphorus are particularly severe: the phosphorus continues to burn as it penetrates tissues, and particles must be removed under water as they reignite in air. Historically, workers in white phosphorus match factories developed a terrible disease called "phossy jaw" (phosphorus necrosis of the jaw), leading to its ban in matches in the early 20th century.
