Phosphorus

What is Phosphorus?

The chemical element phosphorus has the atomic number fifteen and the symbol P. White and red phosphorus are the two main forms of elemental phosphorus; nevertheless, due to its strong reactivity, phosphorus is never found on Earth as a free element.

It is found in the crust of the Earth at a concentration of roughly one gram per kilogram (approximately 0.06 grams for copper).

Phosphorus typically manifests as phosphate in minerals. In 1669, white phosphorus was first extracted and identified as elemental phosphorus.

Phosphorus atoms are grouped in groups of four in white phosphorus, denoted as P₄. The name “white phosphorus” comes from Greek mythology, Φωσφόρoς, which means “light-bearer” (Latin Lucifer) and refers to the planet Venus, sometimes known as the “Morning Star.”

White phosphorus displays a weak glow when exposed to oxygen. This characteristic of phosphorus gives rise to the term “phosphorescence,” which means “glow after illumination,” however it is now also used to refer to a distinct physical process that generates a glow.

Oxidation of white (but not red) phosphorus produces the glow in phosphorus; this phenomenon is known as chemiluminescence.

Phosphorus, along with nitrogen, arsenic, bismuth, antimony, and moscovium, is categorized as a pnictogen.

Phosphorus, which is necessary for life, is mostly obtained from phosphates, which are substances that contain the phosphate ion, PO4³⁻.

Phosphates are part of the complex molecules that make up cells, such as phospholipids, DNA, and RNA.

The first known source of elemental phosphorus was human urine, and an important early source of phosphate was bone ash.

Fossils can be found in phosphate mines because fossilized animal remains and excrement contain phosphate.

Low phosphorus concentrations are a significant growth constraint in many plant habitats. The great majority of fertilizers are made from mined phosphorus compounds.

The yearly demand for phosphate is increasing at a rate that is almost twice as fast as the rate at which plants extract phosphorus from the soil.

Organophosphorus chemicals are also used as nerve agents, insecticides, and detergents.

Characteristics

Allotropes

There are multiple allotropes of phosphorus, and they each have remarkably different properties. White and red phosphorus are the two most prevalent allotropes.

White phosphorus, or WP for short, is the most significant form of elemental phosphorus as seen through the lenses of applications and chemical literature.

It is a soft, waxy solid made up of tetrahedral P₄ molecules, in which a formal single bond binds each of the four atoms to the other three.

Moreover, this P₄ tetrahedron can be found in liquid and gaseous phosphorus up to 800 °C (1,470 °F), at which point it begins to break down into P₂ molecules.

Gas electron diffraction revealed that the P-4 molecule in the gas phase had a P-P bond length of rg = 2.1994(3) Å.

The P₄ tetrahedron’s bonding nature can be explained by either cluster bonding or spherical aromaticity, which indicates that the electrons are substantially delocalized.

Calculations of the magnetically generated currents have demonstrated this; they total 29 nA/T, which is significantly larger than that of the archetypal aromatic molecule benzene (11 nA/T).

There are two crystalline forms of white phosphorus: α (alpha) and β (beta).

The α-form is stable at room temperature. It is more prevalent, has a cubic crystal structure, and changes into the β-form, which has a hexagonal crystal structure, at 195.2 K (−78.0 °C).

The P4 tetrahedra that make up these shapes have different relative orientations. White phosphorus in its β form has three slightly different P-4 molecules, or eighteen different P-P bond lengths ranging from 2.1768 to 2.1920 Å.

P-P bonds have an average length of 2.183 Å. Of all the phosphorus allotropes, white phosphorus is the least dense, least stable, most reactive, volatile, and most poisonous. Red phosphorus progressively transforms from white phosphorus.

Heat and light speed up this process, therefore samples of white phosphorus that appear yellow really almost always contain some red phosphorus.

Because of this, aged or otherwise impure white phosphorus (i.e., weapons-grade WP rather than lab-grade WP) is sometimes known as yellow phosphorus.

White phosphorus glows in the dark with a very subtle blue and green tint when it is exposed to oxygen.

When in contact with air, it becomes extremely flammable and pyrophoric (self-igniting). Because of its White phosphorus, or pyrophoricity, is an ingredient in napalm.

This type burns with a distinct garlic-like smell, and samples are usually coated in white phosphorus pentoxide, which is made up of P₄O₁₀ tetrahedra with oxygen inserted at the vertices and between the phosphorus atoms.

While soluble in carbon disulfide, white phosphorus is insoluble in water. P₄ thermally decomposes to diphosphorus, P₂, at 1100 K.

Either as a liquid or a solid, this species is unstable. The dimeric unit is similar to N2 and has a triple bond.

It can also be produced in solution as a temporary intermediate by thermalizing reagents that are precursors to organophosphorus.

Higher temperatures cause P₂ to split off into atomic P. The structure of red phosphorus is polymeric.

It can be thought of as a derivative of P₄, with chains of P₂₁ molecules bound together by van der Waals forces as a result of one P-P connection being broken and one new bond being created with the nearby tetrahedron.

White phosphorus can be heated to 250 °C (482 °F) or exposed to sunshine to produce red phosphorus.

Following this treatment, phosphorus is amorphous. More heating causes this substance to crystallize.

Red phosphorus is therefore an intermediate phase between the white and blue phosphorus, not an allotrope.

Violet phosphorus and the majority of its characteristics are not fixed. Bright red phosphorus, for instance, ignites at approximately 300 °C (572 °F) when it is freshly manufactured.

It is more stable than white phosphorus, which ignites at approximately 30 °C (86 °F). The color darkens (see infobox photographs) after extended heating or storage.

The resulting product is more stable and doesn’t ignite spontaneously in the air. Red phosphorus can be annealed for a whole day at temperatures exceeding 550 °C to create violet phosphorus.

Hittorf found in 1865 that phosphorus takes on a reddish-purple appearance when it recrystallizes from molten lead.

Hence, this kind is occasionally referred to as “Hittorf’s phosphorus” (also called violet or α-metallic phosphorus).

Below 550 °C (1,022 °F), black phosphorus is the thermodynamically stable form and the least reactive allotrope.

It is also referred to as β-metallic phosphorus and possesses a structure that is a little similar to graphite.

It is produced by heating white phosphorus to a high pressure of 1.2 gigapascals, or 12,000 standard atmospheres.

It can also be made at room temperature with metal salts as catalysts, such as mercury. It is similar to graphite in structure, appearance, and properties.

It is a black, brittle substance that conducts electricity and features puckered sheets of connected atoms.

By letting a solution of white phosphorus in carbon disulfide evaporate in the presence of sunshine, scarlet phosphorus—another form—is produced.

Chemiluminescence

When white phosphorus was originally isolated, it was found that the green glow it produced in a stoppered jar would last for a while before ceasing.

In the 1680s, Robert Boyle attributed it to the “debilitation” of the air. In actuality, oxygen is being used up.

Phosphorus glows exclusively over a range of partial pressures; in pure oxygen, it does not glow at all, as was discovered during the 18th century.

Higher pressures can be achieved by applying heat to fuel the process. A. U. Khan and R. J. van Zee explained luminescence in 1974.

The surface of the solid (or liquid) phosphorus undergoes a reaction with oxygen to generate the short-lived molecules HPO and P₂O₂, which both emit visible light.

The glow in a stoppered jar lasts for a long time because the reaction is slow and just a small amount of the intermediates are needed to produce the luminescence.

.Phosphorus and phosphorescence have been used synonymously to refer to materials that glow in the dark without burning ever since they were discovered.

Phosphorus glows due to a process known as chemiluminescence, which is the correct term for the reaction.

Not phosphorescence (re-emitting light that previously fell onto a substance and aroused it), despite the fact that the term “phosphorescence” is derived from phosphorus.

Isotopes

Phosphorus has 22 known isotopes, which range in value from ²⁶P to ⁴⁷P. Since only ³¹P is stable, it is present in 100% abundance.

Phosphorus-31 NMR spectroscopy is a particularly valuable analytical tool in studies of phosphorus-containing substances because of the half-integer nuclear spin and high abundance of ³¹P.

There are two radioactive isotopes of phosphorus with half-lives that are appropriate for scientific studies on biology. These are the following:

• In life science labs, ³²P is a commonly used beta-emitter (1.71 MeV) with a half-life of 14.3 days. Its main purpose is to create radiolabeled DNA and RNA probes, such as those used in Northern blots and Southern blots.

• ³³P is a beta-emitter with a half-life of 25.4 days and a 0.25 MeV. It is employed in life science labs for tasks like DNA sequencing where reduced energy beta emissions are beneficial.

Skin and corneas are penetrated by the high-energy beta particles from ³²P, and any ³²P that is swallowed, breathed, or absorbed is easily absorbed into bone and nucleic acids.

Because of these factors, workers handling ³²P are required by the Occupational Safety and Health Administration (OSHA).

United States and comparable organizations in other developed nations wear lab coats, disposable gloves, safety glasses, or goggles to protect their eyes, and to avoid working directly over open containers.

It’s also necessary to keep an eye on surface, clothing, and personal contamination. Protecting needs particular attention.

Due to the high energy of the beta particles, dense shielding materials like lead can produce secondary X-ray emission through braking radiation, or Bremsstrahlung.

Thus, low-density materials like wood, water, acrylic, or other plastic—or even wood, if transparency is not needed—must be used to hide the radiation.

Occurrence

Universe

Phosphorus was discovered in Cassiopeia A in 2013, confirming that supernovae produce this element as a byproduct of supernova nucleosynthesis.

Material from the supernova remnant may have a phosphorus-to-iron ratio up to 100 times higher than that of the Milky Way as a whole.

In 2020, scientists examined ALMA and ROSINA data from the AFGL 5142 enormous star-forming area to identify compounds containing phosphorus and their transportation via comets to the early Earth.

Form	white(α)	white(β)	red	violet	black
Symmetry	Body-centered cubic	Triclinic	Amorphous	Monoclinic	Orthorhombic
Pearson symbol		aP24		mP84	oS8
Space group	I43m	P1 No.2		P2/c No.13	Cmce No.64
Density (g/cm3)	1.828	1.88	~2.2	2.36	2.69
Band gap (eV)	2.1		1.8	1.5	0.34
Refractive index	1.8244			2.6	2.4

Crust and Organic Sources

The concentration of phosphorus in the Earth’s crust is around one gram per kilogram (roughly 0.06 grams for copper).

It is widely distributed in numerous minerals, mainly as phosphates, rather than being found freely in the natural world.

The mineral known as apatite, which is primarily composed of pentacalcium, is a component of inorganic phosphate rock.

Nowadays, tri orthophosphate fluoride (hydroxide) is the main source of this element for commerce.

The US Geological Survey (USGS) estimates that Amazigh countries like Morocco, Algeria, and Tunisia contain roughly half of the world’s phosphorus deposits.

Morocco holds 85% of the world’s known reserves, with minor concentrations found in China, Russia, Florida, Idaho, Tennessee, Utah, and other places.

For example, phosphate rock from Tennessee, Florida, and the Îles du Connétable (phosphate island sources) was used by Albright and Wilson in the UK and their Niagara Falls factory in the 1890s and 1900s; by 1950.

However, they were mostly using phosphate rock from Tennessee and North Africa. Although they were historically significant, organic sources such as urine, bone ash, and (in the late 19th century) guano had little economic success.

Urine has fertilizing properties since it includes phosphorus, and some countries, like Sweden, still exploit these features today by repurposing excreta in different ways.

Purine, therefore, can be utilized as a fertilizer either in its undiluted state or in combination with water to create sewage or sewage sludge.

Compounds

Phosphorus(V)

The most common forms of phosphorus are derived from the tetrahedral anion phosphate (PO43−).

The conjugate base of phosphoric acid, which is generated in large quantities for use in fertilizers, is phosphate.

Phosphoric acid, being tricrotic, progressively transforms into three conjugate bases:
H₃PO₄ + H₂O ⇌ H₃O⁺ + H₂PO₄⁻       K_a1 = 7.25×10⁻³
H₂PO₄⁻ + H₂O ⇌ H₃O⁺ + HPO₄²⁻       K_a2 = 6.31×10⁻⁸
HPO₄²⁻ + H₂O ⇌ H₃O⁺ +  PO₄³⁻        K_a3 = 3.98×10⁻¹³

Phosphate has a propensity to create P-O-P bond-containing chains and rings. ATP is one of the several known polyphosphates.

Dehydration of hydrogen phosphates, such as HPO₄²⁻ and H₂PO₄⁻., produces polyphosphate.

For instance, the megatonne produces the industrially significant pentasodium triphosphate (sometimes called sodium tripolyphosphate, or STPP) by this condensation reaction:

2 Na₂HPO₄ + NaH₂PO₄ → Na₅P₃O₁₀ + 2 H₂O

The acid anhydride of phosphoric acid is phosphorus pentoxide (P₄O₁₀), however, there are a number of recognized intermediates between the two.

Water and this waxy white substance react violently. Phosphate reacts with metal cations to generate a range of salts.

These polymeric solids have P-O-M connections. The salts are often insoluble when the metal cation has a charge of 2+ or 3+, which is why they exist as common minerals. Hydrogen phosphate (HPO₄²⁻) is the source of several phosphate salts.

PCl₅ and PF₅ are typical substances. The molecules of PF₅, a colorless gas, are trigonal bipyramidal in shape.

PCl₅ is a colorless solid with the ionic composition PCl₄⁺ PCl₆⁻; nevertheless, when molten or in the vapor phase, it takes on the geometry of a trigonal bipyramid.

Formulated asPBr₄⁺Br⁻, PBr₅ is an unstable solid; PI₅ is unknown. Lewis acids are pentachloride and pentafluoride.

PF₅ produces PF₆⁻, an anion that is isoelectronic with SF₆ when fluoride is added. Phosphorus oxychloride, or POCl₃, is the most significant oxyhalide.

It is roughly tetrahedral in shape. It was believed that d orbitals were involved in bonding in phosphorus(V) compounds prior to the feasibility of comprehensive computer computations.

Molecular orbital theory computer simulation suggests that only s- and p-orbitals are involved in this bonding.

Phosphorus(III)

The four symmetrical trihalides—the solid PI3, the yellowish liquids PCl₃ and PBr₃, and the gaseous PF₃—are all widely recognized.

These materials hydrolyze to produce phosphorous acid and are moisture-sensitive. White phosphorus is chlorinated to create trichloride, a common reagent:
P₄ + 6 Cl₂ → 4 PCl₃
Through halide exchange, the trichloride is converted to trifluoride.

Because PF₃ binds to hemoglobin, it is hazardous.Phosphorus(III) oxide, P₄O₆, is the minor tautomer of phosphorous acid and is also known as tetraphosphorus hexoxide.

Without the terminal oxide groups, the structure of P₄O₆ is similar to that of P₄O₁₀.

Phosphorus(I) and Phosphorus(II)

P-P bonds are typically present in these molecules. Catenated phosphine derivatives and organophosphines are two examples.

Although they are uncommon, compounds with P=P double bonds have also been identified.

Phosphides and Phosphines

Red phosphorus and metals react to form phosphates. Group 1 alkali metals and alkaline earth metals can combine to generate ionic compounds that include P³⁻, the phosphide ion.

Phosphine is created when these substances combine with water. These reactive metals are recognized to be present in other phosphides, such as Na₃P₇.

There are two types of phosphorus-rich phosphides that are less stable and include semiconductors, and metal-rich phosphides that are typically hard refractory compounds with a metallic luster when combined with transition metals.

Meteorites naturally include schreibersite, a phosphide that is rich in metals. Both phosphorus- and metal-rich phosphides can have intricate structures.

Ammonia (NH₃) and phosphine (PH₃) are structurally similar, but phosphine and its chemical derivatives have bond angles at phosphorus that are closer to 90°.

A poisonous, foul-smelling gas is phosphorus. In phosphine, phosphorus has an oxidation number of -3. Calcium phosphide, or Ca₃P₂, is hydrolyzed to create phosphorus.

Air oxidizes phosphine, unlike ammonia. Furthermore, phosphorus is significantly less basic than ammonia.

There are other phosphines with the formula PnHn+2 that feature chains of up to nine phosphorus atoms. Hydrazine’s equivalent is diphosphine (P₂H₄), a highly flammable gas.

Oxoacids

Numerous, frequently significant from a business standpoint, and occasionally structurally intricate are phosphorus oxoacids.

All of them have acidic protons attached to oxygen atoms, whereas some also contain phosphorus–phosphorus connections and nonacidic protons directly connected to phosphorus.

Only nine phosphorus oxoacids are commercially significant, however many are generated; hypophosphorous acid, phosphorous acid, and phosphoric acid are three of the most significant.

Oxidation state	Formula	Name	Acidic protons	Compounds
+1	HH2PO2	hypophosphorous acid	1	acid, salts
+3	H3PO3	phosphorous acid (phosphonic acid)	2	acid, salts
+3	HPO2	metaphosphorous acid	1	salts
+4	H4P2O6	hypophosphoric acid	4	acid, salts
+5	(HPO3)n	metaphosphoric acids	n	salts (n = 3,4,6)
+5	H(HPO3)nOH	polyphosphoric acids	N+2	acids, salts (n = 1-6)
+5	H5P3O10	tripolyphosphoric acid	3	salts
+5	H4P2O7	pyrophosphoric acid	4	acid, salts
+5	H3PO4	(ortho)phosphoric acid	3	acid, salts

Nitrides

Although the PN molecule is thought to be unstable, it is a byproduct of the breakdown of crystalline phosphorus nitride around 1100K.

Phosphorus nitride halogens such as  F₂PN, Cl₂PN, Br₂PN, and I₂PN oligomerize into cyclic polyphosphazene, however, H₂PN is regarded as unstable.

For instance, molecules of the formula (PNCl₂)n, like trimer hexachlorophosphazene, are generally found as rings.

Phosphorous pentachloride is treated with ammonium chloride to produce phosphazenes:
PCl₅ + NH₄Cl → 1/n (NPCl₂)_n + 4 HCl
Alkoxide (RO⁻) replaces the chloride groups, producing a family of polymers with perhaps helpful characteristics.

Sulfides

Phosphorus can be in P(V), P(III), or other oxidation states and creates a large variety of sulfides. Strike-anywhere matches use the three-fold symmetric P₄S₃. P₄S₁₀ and P₄O₁₀ exhibit structural similarities. Phosphorus(III) mixed oxyhalides and oxyhydrides are hardly known.

Organophosphorus Compounds

Organophosphorus compounds are frequently defined as compounds containing P-C and P-O-C linkages.

They are extensively employed in commerce. In the process of creating organophosphorus(III) compounds, PCl₃ acts as a source of P3⁺. It is the precursor, for instance, to triphenylphosphine:

PCl₃ + 6 Na + 3 C₆H₅Cl → P(C₆H₅)₃ + 6 NaCl
Phosphites, such as triphenylphosphite, are produced when phosphorus trihalides are treated with alcohols and phenols.
PCl₃ + 3 C₆H₅OH → P(OC₆H₅)₃ + 3 HCl
Similar reactions occur for phosphorus oxychloride, affording triphenylphosphate:
OPCl₃ + 3 C₆H₅OH → OP(OC₆H₅)₃ + 3 HCl

History

Etymology

Venus was known by the Greek name Phosphorus, which comes from the Greek words φῶς (light) and φέρω (carry), which roughly translate to “light-bringer” or “light-carrier.”

(In Greek mythology and tradition, there are three close homologs: Augerinus (Αυγερινός = morning star, still in use today).

Hesperus or Hesperinus (΄Εσπερoς or Εσπερινός or Αποσπερίτης = evening star, still in use today), and Eosphorus (Εωσφόρος = dawnbearer, not in use for the planet after Christianity).

Phosphorus is the right spelling of the element, according to the Oxford English Dictionary. Since the word “phosphorous” is the adjectival form of the P³⁺ valence.

Phosphorus forms P⁵⁺ valence phosphoric compounds (such as phosphoric acids and phosphates) and phosphorous compounds (such as phosphorous acid), much as sulfur forms sulfurous and sulfuric compounds.

Discovery

The German alchemist Hennig Brand is credited with discovering phosphorus in 1669, but other people may have made similar discoveries at the same time.

Phosphorus is the first element to be discovered that has been unknown since antiquity. The brand conducted an experiment using urine, which has significant levels of dissolved phosphates from regular metabolism.

While working in Hamburg, Brand tried to make the illusive philosopher’s stone by evaporating urine to condense certain salts; in the process, he created a white substance that lit brightly and gleamed in the dark.

Phosphorus mirabilis was the given name (meaning “miraculous bearer of light”).

In the beginning, Brand’s method was leaving pee alone for days until it started to smell bad. He then reduced it to a paste by boiling it, raising the paste’s temperature, and passing the vapors through water in the hopes that they would condense into gold.

Rather, what he got was a waxy, white substance that shone in the dark. Phosphorus was found by Brand.

In particular, Brand created (NH₄)NaHPO₄ or ammonium sodium hydrogen phosphate. The amounts were roughly right—around 1,100 liters [290 US gal] of urine were needed to produce roughly 60 g of phosphorus—but the urine didn’t need to decay beforehand.

Scientists then found that the amount of phosphorus recovered from fresh pee was the same. Initially, Brand attempted to keep the process a secret, but eventually gave Dresden resident Johann Daniel Kraft (de) the recipe for 200 thalers.

With it, Krafft traveled through most of Europe, stopping in England to meet Robert Boyle.

Following the disclosure of the material’s urine-based composition, Johann Kunckel (1630–1703) was able to replicate it in Sweden (1678).

Later, in London (1680), Boyle also succeeded in producing phosphorus, maybe with the help of Ambrose Godfrey-Hanckwitz, his assistant.

Later on, Godfrey turned phosphorus manufacturing into a profitable venture. Boyle claims that Krafft just told him that phosphorous was made from “something that belonged to the body of man” and provided no further details.

This provided Boyle with a crucial hint, enabling him to successfully produce phosphorous and publish the process recipe.

Afterward, he modified Brand’s procedure by introducing sand into the reaction while maintaining urine as the foundation material.

4 NaPO₃ + 2 SiO₂ + 10 C → 2 Na₂SiO₃ + 10 CO + P₄

In 1680, Robert Boyle became the first person to light sulfur-tipped wooden splints—the precursors of modern matches—with phosphorus.

The thirteenth element to be identified was phosphorus. It’s known as “the Devil’s element” because of its propensity to catch fire when left out in the open.

Bone Ash and Guano

After Johan Gottlieb Gahn and Carl Wilhelm Scheele demonstrated in 1769 that calcium phosphate (Ca₃(PO₄) ₂) is found in bones by getting elemental phosphorus from bone ash,

Antoine Lavoisier identified phosphorus as an element in 1777. Up to the 1840s, the main source of phosphorus was bone ash.

The process began with roasting bones and involved distilling the extremely hazardous elemental phosphorus result using fire clay retorts housed in a very hot brick furnace.

Alternatively, crushed bones that had been cleaned and acid-treated might be used to make precipitated phosphates.

The precipitated phosphates could then be heated, and combined with charcoal or crushed coal in an iron pot, and the phosphorus vapor could be extracted using a retort to create white phosphorus.

In a flare stack, flammable gases such as carbon monoxide that were created during the reduction process were burned off.

The mining of tropical island deposits made of bat and bird guano began to supply the world’s phosphate needs in the 1840s (see also Guano Islands Act).

In the second half of the 1800s, they started to play a significant role as a supply of phosphates for fertilizer.

Phosphate Rock

The first phosphate rock was used to produce phosphorus in 1850. James Burgess Readman invented the electric arc furnace in 1888 (patented in 1889).

After that, phosphate rock—which typically contains calcium phosphate—was used to produce phosphorus instead of heating bone ash.

Following the global guano supply’s depletion during the same period, mineral phosphates emerged as the principal source of Production of phosphate fertilizers.

After World War II, phosphate rock output skyrocketed, and it continues to be the world’s main supply of phosphorus and phosphorus compounds to this day.

For additional details on the background and current situation of phosphate mining, refer to the peak phosphorus article.

To create a variety of “superphosphate” fertilizer products, phosphate rock is still used as a feedstock in the fertilizer business after being treated with sulfuric acid.

Incendiaries

The 19th century saw the commercial manufacture of white phosphorus for the match industry. As previously mentioned, this used bone ash as a supply of phosphate.

When the submerged-arc furnace for phosphorus synthesis was created to decrease phosphate rock, the bone-ash method was rendered obsolete.

Phosphorus may now be employed in military weaponry due to increased production made possible by the electric furnace process.

During the First World War, it was utilized as smoke screens, incendiaries, and tracer bullets. A unique incendiary projectile was created to fire against hydrogen-filled Zeppelin aircraft because hydrogen is extremely combustible, over Britain.

Phosphorus-dissolved gasoline-based Molotov cocktails were given to specifically chosen citizens in Britain as part of the British resistance movement for defense during World War II, and phosphorus incendiary bombs were widely employed in combat.

Phosphorous that has burned has terrible consequences on human skin and is difficult to put out.

White phosphorus, which was hazardous because of its toxicity, was a component of early matches.

Its use led to unintentional poisonings, murders, and suicides. (An apocryphal story describes a woman trying to kill her husband by putting white phosphorus in his dinner, which she discovered by the stew emitting glowing steam).

Moreover, match workers who were exposed to the fumes developed “phossy,” severe necrosis of the jaw bones. jaw”.

Following the discovery of a safe method for producing red phosphorus, which has significantly less flammability and toxicity than matchmaking, regulations were passed under the Berne Convention (1906) mandating its use as a safer substitute for matchmaking.

White phosphorous was no longer used in matches due to its toxicity. The location where the “miraculous bearer of light” was first found, Hamburg, was destroyed by the Allies during World War II with phosphorus incendiary bombs.

Production

0.261 billion tons were mined in 2016, and the USGS projected 68 billion tons of world reserves in 2017.

Reserve numbers relate to the quantity assumed recoverable at current market prices. Its annual demand is growing at a rate almost twice as fast as the expansion of the human population, making it essential to modern agriculture.

The creation of It’s possible that phosphorous peaked before 2011, and some scientists believe the supplies will run out before the end of the twenty-first century.

Because phosphorus makes up only 0.1% of the mass of an ordinary rock, there is an abundant but diluted supply of phosphorus on Earth.

Wet Process

Phosphorous-containing materials are mostly used as agricultural fertilizers. Here, where the purity requirements are low, phosphorus is extracted from phosphate rock using a procedure known as the “wet process.”

Phosphoric acid is produced by treating the minerals with sulfuric acid. The several phosphate salts that result from neutralizing phosphoric acid are what makeup fertilizers.

There is no redox reaction involving phosphorus in the wet process. For every ton of phosphoric acid produced, about five tons of waste phosphogypsum are produced.

Globally, 100–280 Mt of phosphogypsum are thought to be produced annually.

Thermal Process

The high purity requirements for phosphorus in medications, detergents, and food items prompted the creation of the thermal process.

Phosphate minerals are transformed into white phosphorus during this process, which can then be refined through distillation.

Phosphate salts are produced by oxidizing white phosphorus to phosphoric acid and then neutralizing it with a base.

The submerged-arc furnace used for the thermal process uses a lot of energy. Currently, the annual production of elemental phosphorus is at 1,000,000 short tons (910,000 t).

Phosphate rock, or calcium phosphate, is primarily mined in Florida and North Africa. It can be burned to 1,200–1,500 °C along with coke and sand (mainly SiO₂) to generate P₄.

Due to its volatility, the P₄ product is easily isolated:

4 Ca₅(PO₄)₃F + 18 SiO₂ + 30 C → 3 P₄ + 30 CO + 18 CaSiO₃ + 2 CaF₂
2 Ca₃(PO₄)₂ + 6 SiO₂ + 10 C → 6 CaSiO₃ + 10 CO + P₄

Due to iron impurities in the mineral precursors, ferrophosphorus, a basic form of Fe₂P, is one of the side products of the thermal process.

The silicate slag can be utilized as a building material. On occasion, fluoride is recovered for use in fluoridating water.

An even more concerning “mud” is one that has a high concentration of white phosphorus. Because white phosphorus production requires a lot of energy, it is done in huge facilities.

The molten white phosphorus is conveyed. There have been a few significant transportation-related accidents.

Historical Routes

In the past, white phosphorus was industrially extracted from bone ash prior to the advent of mineral-based extractions.

This method uses sulfuric acid to change the tricalcium phosphate in bone ash into monocalcium phosphate:

Ca₃(PO₄)₂ + 2 H₂SO₄ → Ca(H₂PO₄)₂ + 2 CaSO₄

Next, monocalcium phosphate is dried to yield the matching metaphosphate:

Ca(H₂PO₄)₂ → Ca(PO₃)₂ + 2 H₂O

Two-thirds of the weight of calcium metaphosphate is produced as white phosphorus when it is burnt to a white heat (~1300 °C) with charcoal; the other third of the phosphorus is left as calcium orthophosphate.

3 Ca(PO₃)₂ + 10 C → Ca₃(PO₄)₂ + 10 CO + P₄

Applications

Flame Retardant

Compounds containing phosphorus are utilized as flame retardants. Phosphorus- and bio-based flame-retardant materials and coatings are under development.

Food Additive

According to the Dietary Reference Intake, phosphorus is a necessary mineral for humans (DRI).

To provide a tangy or sour taste to foods and beverages like jams and colas, food-grade phosphoric acid (additive E338) is utilized.

Additionally, phosphoric acid acts as a preservative. Phosphate sodas or phosphates are other names for soft drinks that include phosphoric acid, such as Coca-Cola.

Soft drink phosphoric acid has the ability to erode teeth. Additionally, phosphoric acid may play a role in the development of kidney stones, particularly in individuals who have already experienced kidney stones.

Fertiliser

As the most frequently limiting component after nitrogen, phosphorus is a critical plant nutrient. Most phosphorus is produced in concentrated phosphoric acids for agricultural fertilizers, which can contain as much as 70% to 75% P₂O₅.

As a result, the production of phosphate (PO₄^3-) increased significantly in the second half of the 20th century.

All living things require phosphorus, which is why artificial phosphate fertilization is required.

Phosphorus is involved in energy transfers, the strength of roots and stems, photosynthesis, the expansion of plant roots, the formation of seeds and flowers, and other critical processes that affect the general health and genetics of plants.

Aquatic habitats have been overenriched, or eutrophicated, as a result of excessive application of phosphorus fertilizers and associated discharge.

Because of their poor solubility and poor mobility in soil, naturally occurring phosphorus-containing compounds are largely unavailable to plants.

The majority of phosphorus is highly stable in the soil’s organic matter or minerals. Phosphorus can be fixed in the soil even after being added to manure or fertilizer.

As a result, the phosphorus cycle occurs naturally very slowly. Over time, some of the fixed phosphorus is released again, supporting the growth of wild plants; intense agricultural production requires more.

Superphosphate of lime, a combination of calcium dihydrogen phosphate (Ca(H₂PO₄)₂), and calcium sulfate dihydrate (CaSO₄·2H₂O), which is created when sulfuric acid and water react with calcium phosphate, are common forms of fertilizer.

The world economy depends so heavily on the processing of phosphate minerals with sulfuric acid to create fertilizer that this is the main industrial market for sulfuric acid and the largest application of elemental sulfur in industry.

Widely used compounds	Use
Ca(H2PO4)2·H2O	Baking powder and fertilizers
CaHPO4·2H2O	Baking powder and fertilizers
H3PO4	Manufacture of phosphate fertilisers
PCl3	Manufacture of phosphate fertilizers
POCl3	Manufacture of plasticiser
P4S10	Manufacture of plasticizer
Na5P3O10	Detergents

Organophosphorus

White phosphorus is extensively utilized in the synthesis of two phosphorus sulfides, phosphorus pentasulfide, and phosphorus sesquisulfide, as well as intermediate phosphorus chlorides to create organophosphorus compounds.

Organophosphorus compounds are used in a wide range of products, such as water treatment, insecticides, plasticizers, flame retardants, and extraction agents.

Metallurgical Aspects

In addition, phosphorus is necessary for the synthesis of steel, phosphor bronze, and several other goods that are linked.

During the smelting process, phosphorus is added to metallic copper to react with oxygen, which is an impurity in copper.

This produces alloys called phosphorus-containing copper (CuOFP), which are more resistant to hydrogen embrittlement than regular copper.

Steel components can be chemically treated to increase their resistance to corrosion by applying phosphate conversion coating.

Matches

Charles Sauria created the first striking match with a phosphorous head in 1830. Heads of white phosphorus, an oxygen-releasing substance (potassium chlorate, lead dioxide, or occasionally nitrate), and a binder were used to create these matches and their ensuing variants.

They were toxic if consumed, hazardous if inadvertently burned on a rough surface, noxious to workers during the manufacturing process, and sensitive to storage conditions.

From 1872 until 1925, production was outlawed in a number of nations. Adopted in 1906, the worldwide Berne Convention forbade the use of white phosphorus in matches.

As a result, safer substitutes for phosphorous matches were progressively introduced. The current strike-anywhere match was created by French chemists Henri Sévène and Emile David Cahen in 1900.

Phosphorus sesquisulfide (P₄S₃), a non-pyrophoric and non-toxic substance that ignites under friction, was used in place of white phosphorus.

These safer strike-anywhere matches gained popularity for a while, but the modern safety match eventually supplanted them.

It is quite difficult to ignite a safety match anywhere other than on a designated striker strip. The match head potassium chlorate, which releases oxygen, and non-toxic red phosphorus are both included in the strip.

Little amounts of abrasion from the striker strip and match head are combined carefully when struck to create a little amount of Armstrong’s mixture, a composition that is extremely sensitive to touch.

The tiny powder instantly ignites and offers the first spark that ignites the match head. Until the match is struck, safety matches keep the two parts of the ignition mixture apart.

The main safety benefit is that it keeps accidental ignition from happening. However, until the banning of white phosphorus, safety matches—which Gustaf Erik Pasch devised in 1844.

And had ready for the market by the 1860s—were not widely accepted by consumers. Employing a certain striker strip was deemed awkward.

Water Softening

Laundry detergents contain sodium tripolyphosphate, which is derived from phosphoric acid and is permitted in certain nations but prohibited in others. This substance softens the water to improve detergent efficacy and stop corrosion in pipes and boiler tubes.

Miscellaneous

• To create customized lenses for sodium lights, phosphates are utilized.
• Fine china is made from bone ash, which is mostly calcium phosphate.
• Made from elemental phosphorus, phosphoric acid is used to make food-grade phosphates and is utilized in beverages like soft drinks. These include sodium tripolyphosphate and monocalcium phosphate for baking powder. Phosphates are added to toothpaste and processed meat and cheese to enhance their quality.
• White phosphorous, often known as “WP” (slang for “Willie Peter”), is utilized in tracer ammunition, smoke-screening devices like pots and bombs, and incendiary bombs for military purposes. It is also a component of a defunct US hand grenade, the M34 White Phosphorus. Mostly employed for signaling, smoke screens, and inflammation, this multipurpose bomb could also inflict serious burns and have a psychological effect on the adversary. International law restricts the use of white phosphorous for military purposes.
• In biochemical laboratories, radioactive tracers 32P and 33P are utilized.

Biological Role

All known life forms require inorganic phosphorus, which is present as the phosphate PO³⁻
₄.

A significant part of DNA and RNA’s structural makeup is phosphorus. Phosphate is used by living cells to transfer adenosine triphosphate (ATP), which is required for all energy-consuming cellular processes.

Phosphorylation is another essential cellular regulation process that depends on ATP. The primary structural elements of all cellular membranes are phospholipids.

Salts containing calcium phosphate help to strengthen bones. The term “Pi” is frequently used by biochemists to refer to inorganic phosphate.

A membrane enveloping each live cell keeps it isolated from its environment. Proteins, usually arranged in a bilayer, and a phospholipid matrix make up cellular membranes.

Glycerol is the starting material for phospholipids; two of the hydroxyl (OH) protons are substituted by fatty acids in the form of an ester, and the third hydroxyl proton is replaced with phosphate that is linked to another alcohol.

About 85–90% of the 0.7 kg of phosphorus that an adult human has is found in their bones and teeth (apatite), with the remaining portion found in their soft tissues and extracellular fluids.

From roughly 0.5% by mass in infancy to 0.65–1.1% by mass in adulthood, the phosphorus level rises.

The blood has an average phosphorus concentration of 0.4 g/L, of which 30% is made up of inorganic phosphates and 70% is organic phosphates.

An adult who follows a nutritious diet eats and excretes 1-3 grams of phosphorus daily, mostly in the form of phosphate ions such as H₂PO⁻₄  and  HPO²⁻₄, with consumption occurring in the.

Form of inorganic phosphate and phosphorus-containing biomolecules such nucleic acids and phospholipids.

Comparable to the quantity of phosphate available to soft tissue cells, only approximately 0.1% of the body’s total phosphate is circulating in the blood.

Bone and Teeth Enamel

Hydroxyapatite and other amorphous calcium phosphate forms, potentially including carbonate, make up the majority of bone.

Tooth enamel is primarily composed of hydroxyapatite. Fluoride in water partially transforms this mineral into the still-harder fluorapatite, strengthening teeth’s resistance to decay:
Ca₅(PO₄)₃OH + F⁻→ Ca₅(PO₄)₃F + OH⁻

Phosphorus Deficiency

Phosphate deficiency syndrome (PDS) in medical terms can be brought on by starvation, impaired phosphate absorption, and metabolic syndromes.

That pulls phosphate from the blood (refeeding syndrome, for example) or excrete excessive amounts of it in the urine.

Hypophosphatemia, a disorder characterized by low amounts of soluble phosphate in the blood serum and inside the cells, affects all of them.

Signs and symptoms of Hypophosphatemia are characterized by ATP deficiency leading to muscle and blood cell breakdown as well as brain impairment.

Overdosing on phosphate can cause diarrhea, cause soft tissues and organs to become calcified, and hinder the body’s absorption of iron, calcium, magnesium, and zinc.

In edaphology, phosphorus is a macromineral that is crucial for plant growth and is thoroughly examined to comprehend plant uptake from soil systems.

Many ecosystems are influenced by phosphorus deficiency, which slows down the rate at which organisms grow.

Overdosing on phosphorus can also be detrimental, particularly in aquatic environments where eutrophication can occasionally result in algal blooms.

Nutrition

Dietary Recommendations

In 1997, the U.S. Institute of Medicine (IOM) revised the Recommended Dietary Allowances (RDAs) and Estimated Average Requirements (EARs) for phosphorus.

An estimate known as Adequate Intake (AI) is used in place of EARs and RDAs when there is insufficient data to establish them.

For individuals 19 years of age and older, the current EAR for phosphorus is 580 mg/day. A 700 mg daily RDA is set.

RDAs are greater than EARs in order to determine the levels that will folks whose needs are greater than usual.

Additionally, the RDA for nursing and pregnancy is 700 mg/day. The recommended daily allowance (RDA) for those aged 1 to 18 rises from 460 to 1250 mg.

When there is enough data to support a claim, the IOM establishes acceptable upper intake levels (ULs) for vitamins and minerals.

The upper limit of tolerance for phosphorus is 4000 mg/day. Dietary Reference Intakes (DRIs) is the term used to refer to the EARs, RDAs, AIs, and ULs taken together.

The definitions of AI and UL are the same as they are in the US. The AI is established at 550 mg/day for those 15 years of age and above, including those who are pregnant or nursing.

AI 440 mg/day for kids aged 4 to 10 and 640 mg/day for kids aged 11 to 17. The AIs here are lower than the RDAs in the US.

Teens are in greater need than adults in both systems. After reviewing the same safety question, the European Food Safety Authority determined that not enough data was available to establish a UL.

When labeling food and dietary supplements in the United States, the amount in a serving is expressed as a percentage of Daily Value (%DV).

100% of the Daily Value for phosphorus was 1000 mg for labeling purposes, however beginning of May 27, 2016, it was changed to 1250 mg to comply with the RDA.

The previous and new adult daily values are listed in a table at

Food sources

Though proteins don’t include phosphorus, the primary food sources for phosphorus are the same as those that do.

For instance, phosphorus is usually also found in milk, meat, and soy products. Generally speaking, if there is enough protein and calcium in the diet, there should be enough phosphorus.

Precautions

Phosphorous organic compounds are a diverse class of materials, many of which are necessary for life and some of which are highly hazardous.

These fluorophosphate esters are some of the strongest known neurotoxins. Due to their toxicity, a large variety of organophosphorus chemicals are employed as nerve agents against human enemies and as herbicides, insecticides, fungicides, and other pesticides.

The majority of inorganic phosphates are vital nutrients that are comparatively harmless. Because the white phosphorus allotrope burns in the air and leaves behind phosphoric acid residue, it poses a serious risk.

“Phossy jaw” is the term for the necrosis of the jaw caused by chronic white phosphorus overdose.

When consumed in excess, white phosphorus is poisonous and can seriously harm the liver. It can also result in “Smoking Stool Syndrome.”

In the past, treating external exposure to elemental phosphorus involved washing the afflicted region with a solution of 2% copper(II) sulfate, which formed innocuous chemicals that were later removed by washing.

As per the latest US Navy publication, “Treatment of Chemical Agent Casualties and Conventional Military Chemical Injuries: FM8-285: Part 2 Conventional Military Chemical Injuries,”.

Certain countries continue to use cupric (copper(II)) sulfate, which has been utilized by U.S. soldiers in the past.

But since copper sulfate is poisonous, its use will end. Copper sulfate can cause intravascular hemolysis, as well as damage to the kidneys and brain.

“Alternatively, the handbook recommends “neutralizing phosphoric acid with a bicarbonate solution, which will allow removal of visible white phosphorus.”

Particles can frequently be identified by their phosphorescence in the dark or by the smoke they release when air strikes them. In dimly lit environments, pieces appear as bright patches.

If the patient’s condition allows for the removal of any white phosphorus (WP) fragments that may be subsequently absorbed and can result in systemic toxicity.

Applying oily-based ointments should wait until you are positive that all WP has been eliminated.

Treat the lesions as thermal burns when the particles have been completely removed.”[Reference required] Oily materials or ointments are not advised until.

The region has been completely cleaned and all white phosphorus has been eliminated since white phosphorus easily combines with oils.

Phosphorus exposure in the workplace can occur through skin contact, ocular contact, ingestion, and inhalation.

The permissible exposure limit (PEL) for phosphorus in the workplace has been set by the Occupational Safety and Health Administration (OSHA) at 0.1 mg/m3 for an 8-hour workday.

An eight-hour workday is considered to have a Recommended Exposure Limit (REL) of 0.1 mg/m3, according to the.

National Institute for Occupational Safety and Health (NIOSH). Phosphorus poses an urgent risk to life and health at 5 mg/m3.

US DEA List I Status

Ephedrine or pseudoephedrine can be effectively reduced to methamphetamine by phosphorus converting elemental iodine to hydroiodic acid.

Because of this, on November 17, 2001, the US Drug Enforcement Administration listed red and white phosphorus as List I precursor compounds under 21 CFR 1310.02.

Handlers of red or white phosphorus are subject to strict regulatory constraints in the United States.

Notes

White phosphorus, or WP, glows when exposed to air. If WP is present in a wound but is hidden by tissue or fluids like blood serum, it won’t light up until it does. To see properly in a dark room, one must have dark-adapted eyes.

FAQ