Phosphorus (updated)

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Phosphorus,  ₁₅P

waxy white (yellow cut), red (granules center left, chunk center right), and violet phosphorus
General properties
Pronunciation	/ˈfɒsfərəs/ ​(FOS-fər-əs)
Appearance	Colorless, waxy white, yellow, scarlet, red, violet, black
Standard atomic weight (Ar, standard)	30.973761998(5)
Abundance
in the Earth's crust	5.2 (taking silicon as 100)
Phosphorus in the periodic table
Hydrogen Helium Lithium Beryllium Boron Carbon Nitrogen Oxygen Fluorine Neon Sodium Magnesium Aluminium Silicon Phosphorus Sulfur Chlorine Argon Potassium Calcium Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Gallium Germanium Arsenic Selenium Bromine Krypton Rubidium Strontium Yttrium Zirconium Niobium Molybdenum Technetium Ruthenium Rhodium Palladium Silver Cadmium Indium Tin Antimony Tellurium Iodine Xenon Caesium Barium Lanthanum Cerium Praseodymium Neodymium Promethium Samarium Europium Gadolinium Terbium Dysprosium Holmium Erbium Thulium Ytterbium Lutetium Hafnium Tantalum Tungsten Rhenium Osmium Iridium Platinum Gold Mercury (element) Thallium Lead Bismuth Polonium Astatine Radon Francium Radium Actinium Thorium Protactinium Uranium Neptunium Plutonium Americium Curium Berkelium Californium Einsteinium Fermium Mendelevium Nobelium Lawrencium Rutherfordium Dubnium Seaborgium Bohrium Hassium Meitnerium Darmstadtium Roentgenium Copernicium Nihonium Flerovium Moscovium Livermorium Tennessine Oganesson N ↑ P ↓ As silicon ← phosphorus → sulfur
Atomic number (Z)	15
Group	group 15 (pnictogens)
Period	period 3
Block	p-block
Element category	reactive nonmetal
Electron configuration	[Ne] 3s2 3p3
Electrons per shell	2, 8, 5
Physical properties
Phase at STP	solid
Melting point	white: 317.3 K ​(44.15 °C, ​111.5 °F) red: ∼860 K (∼590 °C, ∼1090 °F)
Boiling point	white: 553.7 K ​(280.5 °C, ​536.9 °F)
Density (near r.t.)	white: 1.823 g/cm3 red: ≈2.2–2.34 g/cm3 violet: 2.36 g/cm3 black: 2.69 g/cm3
Heat of fusion	white: 0.66 kJ/mol
Heat of vaporisation	white: 51.9 kJ/mol
Molar heat capacity	white: 23.824 J/(mol·K)
Vapor pressure (white) P (Pa) 1 10 100 1 k 10 k 100 k at T (K) 279 307 342 388 453 549
Vapour pressure (red, b.p. 431 °C) 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, −2, −1, +1, +2, +3, +4, +5
Electronegativity	Pauling scale: 2.19
Ionisation energies	1st: 1011.8 kJ/mol 2nd: 1907 kJ/mol 3rd: 2914.1 kJ/mol (more)
Covalent radius	107±3 pm
Van der Waals radius	180 pm
Spectral lines of phosphorus
Other properties
Natural occurrence	primordial
Crystal structure	​body-centred cubic (bcc)
Thermal conductivity	white: 0.236 W/(m·K) black: 12.1 W/(m·K)
Magnetic ordering	white, red, violet, black: diamagnetic
Magnetic susceptibility	−20.8·10−6 cm3/mol (293 K)
Bulk modulus	white: 5 GPa red: 11 GPa
CAS Number	7723-14-0 (red) 12185-10-3 (white)
History
Discovery	Hennig Brand (1669)
Recognized as an element by	Antoine Lavoisier (1777)
Main isotopes of phosphorus
Iso­tope Abun­dance Half-life (t1/2) Decay mode Pro­duct 31P 100% stable 32P trace 14.28 d β− 32S 33P trace 25.3 d β− 33S

Phosphorus is a chemical element with symbol P and atomic number 15. Elemental phosphorus exists in two major forms, white phosphorus and red phosphorus, but because it is highly reactive, phosphorus is never found as a free element on Earth. It has a concentration in the Earth's crust of about one gram per kilogram (compare copper at about 0.06 grams). With few exceptions, minerals containing phosphorus are in the maximally oxidized state as inorganic phosphate rocks.

Elemental phosphorus was first isolated (as white phosphorus) in 1669 and emitted a faint glow when exposed to oxygen – hence the name, taken from Greek mythology, Φωσφόρος meaning "light-bearer" (Latin Lucifer), referring to the "Morning Star", the planet Venus. The term "phosphorescence", meaning glow after illumination, derives from this property of phosphorus, although the word has since been used for a different physical process that produces a glow. The glow of phosphorus is caused by oxidation of the white (but not red) phosphorus — a process now called chemiluminescence. Together with nitrogen, arsenic, antimony, and bismuth, phosphorus is classified as a pnictogen.

Phosphorus is essential for life. Phosphates (compounds containing the phosphate ion, PO₄³⁻) are a component of DNA, RNA, ATP, and phospholipids. Elemental phosphorus was first isolated from human urine, and bone ash was an important early phosphate source. Phosphate mines contain fossils because phosphate is present in the fossilized deposits of animal remains and excreta. Low phosphate levels are an important limit to growth in some aquatic systems. The vast majority of phosphorus compounds mined are consumed as fertilizers. Phosphate is needed to replace the phosphorus that plants remove from the soil, and its annual demand is rising nearly twice as fast as the growth of the human population. Other applications include organophosphorus compounds in detergents, pesticides, and nerve agents.

Characteristics

Allotropes

White phosphorus exposed to air glows in the dark

 

Crystal structure of red phosphorus

 

Crystal structure of black phosphorus

 

Phosphorus has several allotropes that exhibit strikingly different properties. The two most common allotropes are white phosphorus and red phosphorus.

From the perspective of applications and chemical literature, the most important form of elemental phosphorus is white phosphorus, often abbreviated as WP. It is a soft, waxy solid which consists of tetrahedral P
₄ molecules, in which each atom is bound to the other three atoms by a single bond. This P
₄ tetrahedron is also present in liquid and gaseous phosphorus up to the temperature of 800 °C (1,470 °F) when it starts decomposing to P
₂ molecules. White phosphorus exists in two crystalline forms: α (alpha) and β (beta). At room temperature, the α-form is stable, which is more common and it has cubic crystal structure and at 195.2 K (−78.0 °C), it transforms into β-form, which has hexagonal crystal structure. These forms differ in terms of the relative orientations of the constituent P₄ tetrahedra.

White phosphorus is the least stable, the most reactive, the most volatile, the least dense, and the most toxic of the allotropes. White phosphorus gradually changes to red phosphorus. This transformation is accelerated by light and heat, and samples of white phosphorus almost always contain some red phosphorus and accordingly appear yellow. For this reason, white phosphorus that is aged or otherwise impure (e.g., weapons-grade, not lab-grade WP) is also called yellow phosphorus. When exposed to oxygen, white phosphorus glows in the dark with a very faint tinge of green and blue. It is highly flammable and pyrophoric (self-igniting) upon contact with air. Owing to its pyrophoricity, white phosphorus is used as an additive in napalm. The odour of combustion of this form has a characteristic garlic smell, and samples are commonly coated with white "phosphorus pentoxide", which consists of P
₄O
₁₀ tetrahedra with oxygen inserted between the phosphorus atoms and at their vertices. White phosphorus is insoluble in water but soluble in carbon disulfide.

Thermolysis of P₄ at 1100 kelvin gives diphosphorus, P₂. This species is not stable as a solid or liquid. The dimeric unit contains a triple bond and is analogous to N₂. It can also be generated as a transient intermediate in solution by thermolysis of organophosphorus precursor reagents. At still higher temperatures, P₂ dissociates into atomic P.

Red phosphorus is polymeric in structure. It can be viewed as a derivative of P₄ wherein one P-P bond is broken, and one additional bond is formed with the neighboring tetrahedron resulting in a chain-like structure. Red phosphorus may be formed by heating white phosphorus to 250 °C (482 °F) or by exposing white phosphorus to sunlight. Phosphorus after this treatment is amorphous. Upon further heating, this material crystallizes. In this sense, red phosphorus is not an allotrope, but rather an intermediate phase between the white and violet phosphorus, and most of its properties have a range of values. For example, freshly prepared, bright red phosphorus is highly reactive and ignites at about 300 °C (572 °F), though it is more stable than white phosphorus, which ignites at about 30 °C (86 °F). After prolonged heating or storage, the color darkens (see infobox images); the resulting product is more stable and does not spontaneously ignite in air.

Violet phosphorus is a form of phosphorus that can be produced by day-long annealing of red phosphorus above 550 °C. In 1865, Hittorf discovered that when phosphorus was recrystallized from molten lead, a red/purple form is obtained. Therefore, this form is sometimes known as "Hittorf's phosphorus" (or violet or α-metallic phosphorus).

Black phosphorus is the least reactive allotrope and the thermodynamically stable form below 550 °C (1,022 °F). It is also known as β-metallic phosphorus and has a structure somewhat resembling that of graphite. It is obtained by heating white phosphorus under high pressures (about 12,000 standard atmospheres or 1.2 gigapascals). It can also be produced at ambient conditions using metal salts, e.g. mercury, as catalysts. In appearance, properties, and structure, it resembles graphite, being black and flaky, a conductor of electricity, and has puckered sheets of linked atoms.

Another form, scarlet phosphorus, is obtained by allowing a solution of white phosphorus in carbon disulfide to evaporate in sunlight.

Properties of some allotropes of phosphorus

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

Chemiluminescence

When first isolated, it was observed that the green glow emanating from white phosphorus would persist for a time in a stoppered jar, but then cease. Robert Boyle in the 1680s ascribed it to "debilitation" of the air. Actually, 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 pressures at which it does. 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 the short-lived molecules HPO and P
₂O
₂ that both emit visible light. The reaction is slow and only very little of the intermediates are required to produce the luminescence, hence the extended time the glow continues in a stoppered jar.

Since its discovery, phosphors and phosphorescence were used loosely to describe substances that shine in the dark without burning. Although the term phosphorescence is derived from phosphorus, the reaction that gives phosphorus its glow is properly called chemiluminescence (glowing due to a cold chemical reaction), not phosphorescence (re-emitting light that previously fell onto a substance and excited it).

Isotopes

Twenty-three isotopes of phosphorus are known, including all possibilities from ²⁴P up to ⁴⁶P. Only ³¹P is stable and is therefore present at 100% abundance. The half-integer nuclear spin and high abundance of ³¹P make phosphorus-31 NMR spectroscopy a very useful analytical tool in studies of phosphorus-containing samples. 

Two radioactive isotopes of phosphorus have half-lives suitable for biological scientific experiments. These are:

1. 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.
2. 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.

The high energy beta particles from ³²P penetrate skin and corneas and any ³²P ingested, inhaled, or absorbed is readily incorporated into bone and nucleic acids. For these reasons, Occupational Safety and Health Administration in the United States, and similar institutions in other developed countries require personnel working with ³²P to wear lab coats, disposable gloves, and safety glasses or goggles to protect the eyes, and avoid working directly over open containers. Monitoring personal, clothing, and surface contamination is also required. Shielding requires special consideration. The high energy of the beta particles gives rise to secondary emission of X-rays via Bremsstrahlung (braking radiation) in dense shielding materials such as lead. Therefore, the radiation must be shielded with low density materials such as acrylic or other plastic, water, or (when transparency is not required), even wood.

Occurrence

Universe

In 2013, astronomers detected phosphorus in Cassiopeia A, which confirmed that this element is produced in supernovae as a byproduct of supernova nucleosynthesis. The phosphorus-to-iron ratio in material from the supernova remnant could be up to 100 times higher than in the Milky Way in general.

Crust and organic sources

Phosphorus has a concentration in the Earth's crust of about one gram per kilogram (compare copper at about 0.06 grams). It is not found free in nature, but is widely distributed in many minerals, usually as phosphates. Inorganic phosphate rock, which is partially made of apatite (a group of minerals being, generally, pentacalcium triorthophosphate fluoride (hydroxide)), is today the chief commercial source of this element. According to the US Geological Survey (USGS), about 50 percent of the global phosphorus reserves are in the Arab nations. Large deposits of apatite are located in China, Russia, Morocco, Florida, Idaho, Tennessee, Utah, and elsewhere. Albright and Wilson in the UK and their Niagara Falls plant, for instance, were using phosphate rock in the 1890s and 1900s from Tennessee, Florida, and the Îles du Connétable (guano island sources of phosphate); by 1950, they were using phosphate rock mainly from Tennessee and North Africa.

Organic sources, namely urine, bone ash and (in the latter 19th century) guano, were historically of importance but had only limited commercial success. As urine contains phosphorus, it has fertilising qualities which are still harnessed today in some countries, including Sweden, using methods for reuse of excreta. To this end, urine can be used as a fertiliser in its pure form or part of being mixed with water in the form of sewage or sewage sludge.

Compounds

Phosphorus(V)

The tetrahedral structure of P₄O₁₀ and P₄S₁₀

 

The most prevalent compounds of phosphorus are derivatives of phosphate (PO₄³⁻), a tetrahedral anion. Phosphate is the conjugate base of phosphoric acid, which is produced on a massive scale for use in fertilizers. Being triprotic, phosphoric acid converts step-wise to 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 exhibits a tendency to form chains and rings containing P-O-P bonds. Many polyphosphates are known, including ATP. Polyphosphates arise by dehydration of hydrogen phosphates such as HPO₄²⁻ and H₂PO₄⁻. For example, the industrially important pentasodium triphosphate (also known as sodium tripolyphosphate, STPP) is produced industrially on by the megatonne by this condensation reaction:

2 Na₂[(HO)PO₃] + Na[(HO)₂PO₂] → Na₅[O₃P-O-P(O)₂-O-PO₃] + 2 H₂O

Phosphorus pentoxide (P₄O₁₀) is the acid anhydride of phosphoric acid, but several intermediates between the two are known. This waxy white solid reacts vigorously with water.

With metal cations, phosphate forms a variety of salts. These solids are polymeric, featuring P-O-M linkages. When the metal cation has a charge of 2+ or 3+, the salts are generally insoluble, hence they exist as common minerals. Many phosphate salts are derived from hydrogen phosphate (HPO₄²⁻). 

PCl₅ and PF₅ are common compounds. PF₅ is a colorless gas and the molecules have trigonal bipyramidal geometry. PCl₅ is a colorless solid which has an ionic formulation of PCl₄⁺ PCl₆⁻, but adopts the trigonal bipyramidal geometry when molten or in the vapor phase. PBr₅ is an unstable solid formulated as PBr₄⁺Br⁻and PI₅ is not known. The pentachloride and pentafluoride are Lewis acids. With fluoride, PF₅ forms PF₆⁻, an anion that is isoelectronic with SF₆. The most important oxyhalide is phosphorus oxychloride, (POCl₃), which is approximately tetrahedral. 

Before extensive computer calculations were feasible, it was thought that bonding in phosphorus(V) compounds involved d orbitals. Computer modeling of molecular orbital theory indicates that this bonding involves only s- and p-orbitals.

Phosphorus(III)

All four symmetrical trihalides are well known: gaseous PF₃, the yellowish liquids PCl₃ and PBr₃, and the solid PI₃. These materials are moisture sensitive, hydrolyzing to give phosphorous acid. The trichloride, a common reagent, is produced by chlorination of white phosphorus:

P₄ + 6 Cl₂ → 4 PCl₃

The trifluoride is produced from the trichloride by halide exchange. PF₃ is toxic because it binds to haemoglobin. 

Phosphorus(III) oxide, P₄O₆ (also called tetraphosphorus hexoxide) is the anhydride of P(OH)₃, the minor tautomer of phosphorous acid. The structure of P₄O₆ is like that of P₄O₁₀ without the terminal oxide groups.

Phosphorus(I) and phosphorus(II)

A stable diphosphene, a derivative of phosphorus(I).

 

These compounds generally feature P-P bonds. Examples include catenated derivatives of phosphine and organophosphines. Compounds containing P=P double bonds have also been observed, although they are rare.

Phosphides and phosphines

Phosphides arise by reaction of metals with red phosphorus. The alkali metals (group 1) and alkaline earth metals can form ionic compounds containing the phosphide ion, P³⁻. These compounds react with water to form phosphine. Other phosphides, for example Na₃P₇, are known for these reactive metals. With the transition metals as well as the monophosphides there are metal-rich phosphides, which are generally hard refractory compounds with a metallic luster, and phosphorus-rich phosphides which are less stable and include semiconductors. Schreibersite is a naturally occurring metal-rich phosphide found in meteorites. The structures of the metal-rich and phosphorus-rich phosphides can be complex. 

Phosphine (PH₃) and its organic derivatives (PR₃) are structural analogues of ammonia (NH₃), but the bond angles at phosphorus are closer to 90° for phosphine and its organic derivatives. It is an ill-smelling, toxic compound. Phosphorus has an oxidation number of -3 in phosphine. Phosphine is produced by hydrolysis of calcium phosphide, Ca₃P₂. Unlike ammonia, phosphine is oxidized by air. Phosphine is also far less basic than ammonia. Other phosphines are known which contain chains of up to nine phosphorus atoms and have the formula P_nH_n+2. The highly flammable gas diphosphine (P₂H₄) is an analogue of hydrazine.

Oxoacids

Phosphorous oxoacids are extensive, often commercially important, and sometimes structurally complicated. They all have acidic protons bound to oxygen atoms, some have nonacidic protons that are bonded directly to phosphorus and some contain phosphorus - phosphorus bonds. Although many oxoacids of phosphorus are formed, only nine are commercially important, and three of them, hypophosphorous acid, phosphorous acid, and phosphoric acid, are particularly important.

Oxidation state	Formula	Name	Acidic protons	Compounds
+1	HH2PO2	hypophosphorous acid	1	acid, salts
+3	H2HPO3	phosphorous acid	2	acid, salts
+3	HPO2	metaphosphorous acid	1	salts
+3	H3PO3	(ortho)phosphorous acid	3	acid, 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

The PN molecule is considered unstable, but is a product of crystalline phosphorus nitride decomposition at 1100 K. Similarly, H₂PN is considered unstable, and phosphorus nitride halogens like F₂PN, Cl₂PN, Br₂PN, and I₂PN oligomerize into cyclic Polyphosphazenes. For example, compounds of the formula (PNCl₂)_n exist mainly as rings such as the trimer hexachlorophosphazene. The phosphazenes arise by treatment of phosphorus pentachloride with ammonium chloride:

PCl₅ + NH₄Cl → 1/n (NPCl₂)_n + 4 HCl

When the chloride groups are replaced by alkoxide (RO⁻), a family of polymers is produced with potentially useful properties.

Sulfides

Phosphorus forms a wide range of sulfides, where the phosphorus can be in P(V), P(III) or other oxidation states. The most famous is the three-fold symmetric P₄S₃ which is used in strike-anywhere matches. P₄S₁₀ and P₄O₁₀ have analogous structures. Mixed oxyhalides and oxyhydrides of phosphorus(III) are almost unknown.

Organophosphorus compounds

Compounds with P-C and P-O-C bonds are often classified as organophosphorus compounds. They are widely used commercially. The PCl₃ serves as a source of P³⁺ in routes to organophosphorus(III) compounds. For example, it is the precursor to triphenylphosphine:

PCl₃ + 6 Na + 3 C₆H₅Cl → P(C₆H₅)₃ + 6 NaCl

Treatment of phosphorus trihalides with alcohols and phenols gives phosphites, e.g. triphenylphosphite:

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

The name Phosphorus in Ancient Greece was the name for the planet Venus and is derived from the Greek words (φῶς = light, φέρω = carry), which roughly translates as light-bringer or light carrier. (In Greek mythology and tradition, Augerinus (Αυγερινός = morning star, still in use today), Hesperus or Hesperinus (΄Εσπερος or Εσπερινός or Αποσπερίτης = evening star, still in use today) and Eosphorus (Εωσφόρος = dawnbearer, not in use for the planet after Christianity) are close homologues, and also associated with Phosphorus-the-planet). 

According to the Oxford English Dictionary, the correct spelling of the element is phosphorus. The word phosphorous is the adjectival form of the P³⁺ valence: so, just as sulfur forms sulfurous and sulfuric compounds, phosphorus forms phosphorous compounds (e.g., phosphorous acid) and P⁵⁺ valence phosphoric compounds (e.g., phosphoric acids and phosphates).

Discovery

Robert Boyle

 

The discovery of phosphorus, the first element to be discovered that was not known since ancient times, is credited to the German alchemist Hennig Brand in 1669, although other chemists might have discovered phosphorus around the same time. Brand experimented with urine, which contains considerable quantities of dissolved phosphates from normal metabolism. Working in Hamburg, Brand attempted to create the fabled philosopher's stone through the distillation of some salts by evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. It was named phosphorus mirabilis ("miraculous bearer of light").

Brand's process originally involved letting urine stand for days until it gave off a terrible smell. Then he boiled it down to a paste, heated this paste to a high temperature, and led the vapor through water, where he hoped they would condense to gold. Instead, he obtained a white, waxy substance that glowed in the dark. Brand had discovered phosphorus. We now know that Brand produced ammonium sodium hydrogen phosphate, (NH
₄)NaHPO
₄. While the quantities were essentially correct (it took about 1,100 litres [290 US gal] of urine to make about 60 g of phosphorus), it was unnecessary to allow the urine to rot first. Later scientists discovered that fresh urine yielded the same amount of phosphorus.

Brand at first tried to keep the method secret, but later sold the recipe for 200 thalers to D. Krafft from Dresden, who could now make it as well, and toured much of Europe with it, including England, where he met with Robert Boyle. The secret that it was made from urine leaked out and first Johann Kunckel (1630–1703) in Sweden (1678) and later Boyle in London (1680) also managed to make phosphorus, possibly with the aid of his assistant, Ambrose Godfrey-Hanckwitz, who later made a business of the manufacture of phosphorus.

Boyle states that Krafft gave him no information as to the preparation of phosphorus other than that it was derived from "somewhat that belonged to the body of man". This gave Boyle a valuable clue, so that he, too, managed to make phosphorus, and published the method of its manufacture. Later he improved Brand's process by using sand in the reaction (still using urine as base material),

4 NaPO
₃ + 2 SiO
₂ + 10 C → 2 Na
₂SiO
₃ + 10 CO + P
₄

Robert Boyle was the first to use phosphorus to ignite sulfur-tipped wooden splints, forerunners of our modern matches, in 1680.

Phosphorus was the 13th element to be discovered. For this reason, and due to its use in explosives, poisons and nerve agents, it is sometimes referred to as "the Devil's element".

Bone ash and guano

Guano mining in the Central Chincha Islands, ca. 1860.

 

In 1769, Johan Gottlieb Gahn and Carl Wilhelm Scheele showed that calcium phosphate (Ca
₃(PO
₄)
₂) is found in bones, and they obtained elemental phosphorus from bone ash. Antoine Lavoisier recognized phosphorus as an element in 1777. Bone ash was the major source of phosphorus until the 1840s. The method started by roasting bones, then employed the use of clay retorts encased in a very hot brick furnace to distill out the highly toxic elemental phosphorus product. Alternately, precipitated phosphates could be made from ground-up bones that had been de-greased and treated with strong acids. White phosphorus could then be made by heating the precipitated phosphates, mixed with ground coal or charcoal in an iron pot, and distilling off phosphorus vapor in a retort. Carbon monoxide and other flammable gases produced during the reduction process were burnt off in a flare stack. 

In the 1840s, world phosphate production turned to the mining of tropical island deposits formed from bird and bat guano (see also Guano Islands Act). These became an important source of phosphates for fertilizer in the latter half of the 19th century.

Phosphate rock

Phosphate rock, which usually contains calcium phosphate, was first used in 1850 to make phosphorus, and following the introduction of the electric arc furnace by James Burgess Readman in 1888 (patented 1889), elemental phosphorus production switched from the bone-ash heating, to electric arc production from phosphate rock. After the depletion of world guano sources about the same time, mineral phosphates became the major source of phosphate fertilizer production. Phosphate rock production greatly increased after World War II, and remains the primary global source of phosphorus and phosphorus chemicals today. See the article on peak phosphorus for more information on the history and present state of phosphate mining. Phosphate rock remains a feedstock in the fertilizer industry, where it is treated with sulfuric acid to produce various "superphosphate" fertilizer products.

Incendiaries

White phosphorus was first made commercially in the 19th century for the match industry. This used bone ash for a phosphate source, as described above. The bone-ash process became obsolete when the submerged-arc furnace for phosphorus production was introduced to reduce phosphate rock. The electric furnace method allowed production to increase to the point where 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 being highly flammable). During World War II, Molotov cocktails made of phosphorus dissolved in petrol were distributed in Britain to specially selected civilians within the British resistance operation, for defense; 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.

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's giving off luminous steam). In addition, exposure to the vapors gave match workers a severe 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 the Berne Convention (1906), requiring its adoption as a safer alternative for match manufacture. The toxicity of white phosphorus led to discontinuation of its use in matches. The Allies used phosphorus incendiary bombs in World War II to destroy Hamburg, the place where the "miraculous bearer of light" was first discovered.

Production

Mining of phosphate rock in Nauru

 

Most production of phosphorus-bearing material is for agriculture fertilisers. For this purpose, phosphate minerals are converted to phosphoric acid. It follows two distinct chemical routes, the main one being treatment of phosphate minerals with sulfuric acid. The other process utilizes white phosphorus, which may be produced by reaction and distillation from very low grade phosphate sources. The white phosphorus is then oxidized to phosphoric acid and subsequently neutralized with base to give phosphate salts. Phosphoric acid produced from white phosphorus is relatively pure and is the main route for the production of phosphates for all purposes, including detergent production. 

In the early 1990s, Albright and Wilson's purified wet phosphoric acid business was being adversely affected by phosphate rock sales by China and the entry of their long-standing Moroccan phosphate suppliers into the purified wet phosphoric acid business.

Peak phosphorus

In 2017, the USGS estimated 68 billion tons of world reserves, where reserve figures refer to the amount assumed recoverable at current market prices; 0.261 billion tons were mined in 2016. Critical to contemporary agriculture, its annual demand is rising nearly twice as fast as the growth of the human population.

The production of phosphorus may have peaked already (as per 2011), leading to the possibility of global shortages by 2040. In 2007, at the rate of consumption, the supply of phosphorus was estimated to run out in 345 years. However, some scientists now believe that a "peak phosphorus" will occur in 30 years and that "At current rates, reserves will be depleted in the next 50 to 100 years." Cofounder of Boston-based investment firm and environmental foundation Jeremy Grantham wrote in Nature in November 2012 that consumption of the element "must be drastically reduced in the next 20-40 years or we will begin to starve." According to N.N. Greenwood and A. Earnshaw, authors of the textbook, Chemistry of the Elements, however, phosphorus comprises about 0.1% by mass of the average rock, and consequently the Earth's supply is vast, although dilute.

Elemental phosphorus

Presently, about 1,000,000 short tons (910,000 t) of elemental phosphorus is produced annually. Calcium phosphate (phosphate rock), mostly mined in Florida and North Africa, can be heated to 1,200–1,500 °C with sand, which is mostly SiO
₂, and coke (refined coal) to produce vaporized P
₄. The product is subsequently condensed into a white powder under water to prevent oxidation by air. Even under water, white phosphorus is slowly converted to the more stable red phosphorus allotrope. The chemical equation for this process when starting with fluoroapatite, a common phosphate mineral, is:

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

Side products from this process include ferrophosphorus, a crude form of Fe₂P, resulting from iron impurities in the mineral precursors. The silicate slag is a useful construction material. The fluoride is sometimes recovered for use in water fluoridation. More problematic is a "mud" containing significant amounts of white phosphorus. Production of white phosphorus is conducted in large facilities in part because it is energy intensive. The white phosphorus is transported in molten form. Some major accidents have occurred during transportation; train derailments at Brownston, Nebraska and Miamisburg, Ohio led to large fires. The worst incident in recent times was an environmental contamination in 1968 when the sea was polluted from spillage and/or inadequately treated sewage from a white phosphorus plant at Placentia Bay, Newfoundland.

Another process by which elemental phosphorus is extracted includes applying at high temperatures (1500 °C):

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

Historically, before the development of mineral-based extractions, white phosphorus was isolated on an industrial scale from bone ash. In this process, the tricalcium phosphate in bone ash is converted to monocalcium phosphate with sulfuric acid:

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

Monocalcium phosphate is then dehydrated to the corresponding metaphosphate:

Ca(H₂PO₄)₂ → Ca(PO₃)₂ + 2 H₂O

When ignited to a white heat with charcoal, calcium metaphosphate yields two-thirds of its weight of white phosphorus while one-third of the phosphorus remains in the residue as calcium orthophosphate:

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

Applications

Fertilizer

Phosphorus is an essential plant nutrient (often the limiting nutrient), and the bulk of all phosphorus production is in concentrated phosphoric acids for agriculture fertilizers, containing as much as 70% to 75% P₂O₅. Its annual demand is rising nearly twice as fast as the growth of the human population. That led to large increase in phosphate (PO₄³⁻) production in the second half of the 20th century. Artificial phosphate fertilization is necessary because phosphorus is essential to all living organisms; natural phosphorus-bearing compounds are mostly insoluble and inaccessible to plants, and the natural cycle of phosphorus is very slow. Fertilizer is often in the form of superphosphate of lime, a mixture of calcium dihydrogen phosphate (Ca(H₂PO₄)₂), and calcium sulfate dihydrate (CaSO₄·2H₂O) produced reacting sulfuric acid and water with calcium phosphate. 

Processing phosphate minerals with sulfuric acid for obtaining fertiliser is so important to the global economy that this is the primary industrial market for sulfuric acid and the greatest industrial use of elemental sulfur.

Widely used compounds	Use
Ca(H2PO4)2·H2O	Baking powder and fertilizers
CaHPO4·2H2O	Animal food additive, toothpowder
H3PO4	zManufacture of phosphate fertilizers
PCl3	Manufacture of POCl3 and pesticides
POCl3	Manufacture of plasticizer
P4S10	Manufacturing of additives and pesticides
Na5P3O10	Detergents

Organophosphorus

White phosphorus is widely used to make organophosphorus compounds through intermediate phosphorus chlorides and two phosphorus sulfides, phosphorus pentasulfide and phosphorus sesquisulfide. Organophosphorus compounds have many applications, including in plasticisers, flame retardants, pesticides, extraction agents, nerve agents and water treatment.

Metallurgical aspects

Phosphorus is also an important component in steel production, in the making of phosphor bronze, and in many other related products. Phosphorus is added to metallic copper during its smelting process to react with oxygen present as an impurity in copper and to produce phosphorus-containing copper (CuOFP) alloys with a higher hydrogen embrittlement resistance than normal copper.

Matches

Match striking surface made of a mixture of red phosphorus, glue and ground glass. The glass powder is used to increase the friction.

The first striking match with a phosphorus head was invented by Charles Sauria in 1830. These matches (and subsequent modifications) were made with heads of white phosphorus, an oxygen-releasing compound (potassium chlorate, lead dioxide, or sometimes nitrate), and a binder. They were poisonous to the workers in manufacture, sensitive to storage conditions, toxic if ingested, and hazardous when accidentally ignited on a rough surface. Production in several countries was banned between 1872 and 1925. The international Berne Convention, ratified in 1906, prohibited the use of white phosphorus in matches.

In consequence, the 'strike-anywhere' matches were gradually replaced by 'safety matches', wherein the white phosphorus was replaced by phosphorus sesquisulfide (P₄S₃), sulfur, or antimony sulfide. Such matches are difficult to ignite on any surface other than a special strip. The strip contains red phosphorus that heats up upon striking, reacts with the oxygen-releasing compound in the head, and ignites the flammable material of the head.

Water softening

Sodium tripolyphosphate made from phosphoric acid is used in laundry detergents in some countries, but banned for this use in others. This compound softens the water to enhance the performance of the detergents and to prevent pipe/boiler tube corrosion.

Miscellaneous

• Phosphates are used to make special glasses for sodium lamps.
• Bone-ash, calcium phosphate, is used in the production of fine china.
• Phosphoric acid made from elemental phosphorus is used in food applications such as soft drinks, and as a starting point for food grade phosphates. These include mono-calcium phosphate for baking powder and sodium tripolyphosphate. Phosphates are used to improve the characteristics of processed meat and cheese, and in toothpaste.
• White phosphorus, called "WP" (slang term "Willie Peter") is used in military applications as incendiary bombs, for smoke-screening as smoke pots and smoke bombs, and in tracer ammunition. It is also a part of an obsolete M34 White Phosphorus US hand grenade. This multipurpose grenade was mostly used for signaling, smoke screens, and inflammation; it could also cause severe burns and had a psychological impact on the enemy. Military uses of white phosphorus are constrained by international law.
• In trace amounts, phosphorus is used as a dopant for n-type semiconductors.
• ³²P and ³³P are used as radioactive tracers in biochemical laboratories.
• Phosphate is a strong complexing agent for the hexavalent uranyl (UO₂²⁺) species; for this reason, apatite and other natural phosphates can be very rich in uranium.
• Tributylphosphate is an organophosphate soluble in kerosene used to extract uranium in the Purex process for reprocessing spent nuclear fuel.

Biological role

Inorganic phosphorus in the form of the phosphate PO³⁻
₄ is required for all known forms of life. Phosphorus plays a major role in the structural framework of DNA and RNA. Living cells use phosphate to transport cellular energy with adenosine triphosphate (ATP), necessary for every cellular process that uses energy. 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 assist in stiffening bones. Biochemists commonly use the abbreviation "Pi" to refer to inorganic phosphate.

Every living cell is encased in a membrane that separates it from its surroundings. Cellular membranes are composed of a phospholipid matrix and proteins, typically in the form of a bilayer. Phospholipids are derived from glycerol with two of the glycerol hydroxyl (OH) protons replaced by fatty acids as an ester, and the third hydroxyl proton has been replaced with phosphate bonded to another alcohol.

An average adult human contains about 0.7 kg of phosphorus, about 85–90% in bones and teeth in the form of apatite, and the remainder in soft tissues and extracellular fluids (~1%). The phosphorus content increases from about 0.5 weight% in infancy to 0.65–1.1 weight% in adults. Average phosphorus concentration in the blood is about 0.4 g/L, about 70% of that is organic and 30% inorganic phosphates. An adult with healthy diet consumes and excretes about 1–3 grams of phosphorus per day, with consumption in the form of inorganic phosphate and phosphorus-containing biomolecules such as nucleic acids and phospholipids; and excretion almost exclusively in the form of phosphate ions such as H
₂PO⁻
₄ and HPO²⁻
₄. Only about 0.1% of body phosphate circulates in the blood, paralleling the amount of phosphate available to soft tissue cells.

Bone and teeth enamel

The main component of bone is hydroxyapatite as well as amorphous forms of calcium phosphate, possibly including carbonate. Hydroxyapatite is the main component of tooth enamel. Water fluoridation enhances the resistance of teeth to decay by the partial conversion of this mineral to the still harder material called fluoroapatite:

Ca
₅(PO
₄)
₃OH + F⁻ → Ca
₅(PO
₄)
₃F + OH⁻

Phosphorus deficiency

In medicine, phosphate deficiency syndrome may be caused by malnutrition, by failure to absorb phosphate, and by metabolic syndromes that draw phosphate from the blood (such as in refeeding syndrome after malnutrition) or pass too much of it into the urine. All are characterized by hypophosphatemia, which is a condition of low levels of soluble phosphate levels in the blood serum and inside the cells. Symptoms of hypophosphatemia include neurological dysfunction and disruption of muscle and blood cells due to lack of ATP. Too much phosphate can lead to diarrhoea and calcification (hardening) of organs and soft tissue, and can interfere with the body's ability to use iron, calcium, magnesium, and zinc.

Phosphorus is an essential macromineral for plants, which is studied extensively in edaphology to understand plant uptake from soil systems. Phosphorus is a limiting factor in many ecosystems; that is, the scarcity of phosphorus limits the rate of organism growth. An excess of phosphorus can also be problematic, especially in aquatic systems where eutrophication sometimes leads to algal blooms.

Dietary recommendations

The U.S. Institute of Medicine (IOM) updated Estimated Average Requirements (EARs) and Recommended Dietary Allowances (RDAs) for phosphorus in 1997. If there is not sufficient information to establish EARs and RDAs, an estimate designated Adequate Intake (AI) is used instead. The current EAR for phosphorus for people ages 19 and up is 580 mg/day. The RDA is 700 mg/day. RDAs are higher than EARs so as to identify amounts that will cover people with higher than average requirements. RDA for pregnancy and lactation are also 700 mg/day. For children ages 1–18 years the RDA increases with age from 460 to 1250 mg/day. As for safety, the IOM sets Tolerable upper intake levels (ULs) for vitamins and minerals when evidence is sufficient. In the case of phosphorus the UL is 4000 mg/day. Collectively the EARs, RDAs, AIs and ULs are referred to as Dietary Reference Intakes (DRIs).

The European Food Safety Authority (EFSA) refers to the collective set of information as Dietary Reference Values, with Population Reference Intake (PRI) instead of RDA, and Average Requirement instead of EAR. AI and UL defined the same as in United States. For people ages 15 and older, including pregnancy and lactation, the AI is set at 550 mg/day. For children ages 4–10 years the AI is 440 mg/day, for ages 11–17 640 mg/day. These AIs are lower than the U.S RDAs. In both systems, teenagers need more than adults. The European Food Safety Authority reviewed the same safety question and decided that there was not sufficient information to set a UL.

For U.S. food and dietary supplement labeling purposes the amount in a serving is expressed as a percent of Daily Value (%DV). For phosphorus labeling purposes 100% of the Daily Value was 1000 mg, but as of May 27, 2016 it was revised to 1250 mg to bring it into agreement with the RDA. A table of the old and new adult Daily Values is provided at Reference Daily Intake. The original deadline to be in compliance was July 28, 2018, but on September 29, 2017 the FDA released a proposed rule that extended the deadline to January 1, 2020 for large companies and January 1, 2021 for small companies.

Food sources

The main food sources for phosphorus are the same as those containing protein, although proteins do not contain phosphorus. For example, milk, meat, and soya typically also have phosphorus. As a rule, if a diet has sufficient protein and calcium, the amount of phosphorus is probably sufficient.

Precautions

Phosphorus explosion

 

Organic compounds of phosphorus form a wide class of materials; many are required for life, but some are extremely toxic. Fluorophosphate esters are among the most potent neurotoxins known. A wide range of organophosphorus compounds are used for their toxicity as pesticides (herbicides, insecticides, fungicides, etc.) and weaponized as nerve agents against enemy humans. Most inorganic phosphates are relatively nontoxic and essential nutrients.

The white phosphorus allotrope presents a significant hazard because it ignites in air and produces phosphoric acid residue. Chronic white phosphorus poisoning leads to necrosis of the jaw called "phossy jaw". White phosphorus is toxic, causing severe liver damage on ingestion and may cause a condition known as "Smoking Stool Syndrome".

In the past, external exposure to elemental phosphorus was treated by washing the affected area with 2% copper sulfate solution to form harmless compounds that are then 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(II)) 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 white phosphorus. 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. Promptly debride the burn if the patient's condition will permit removal of bits of WP (white phosphorus) that 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 not recommended until the area is thoroughly cleaned and all white phosphorus removed. 

People can be exposed to phosphorus in the workplace by inhalation, ingestion, skin contact, and eye contact. The Occupational Safety and Health Administration (OSHA) has set the phosphorus exposure limit (Permissible exposure limit) in the workplace at 0.1 mg/m³ over an 8-hour workday. The National Institute for Occupational Safety and Health (NIOSH) has set a Recommended exposure limit (REL) of 0.1 mg/m³ over an 8-hour workday. At levels of 5 mg/m³, phosphorus is immediately dangerous to life and health.

US DEA List I status

Phosphorus can reduce elemental iodine to hydroiodic acid, which is a reagent effective for reducing ephedrine or pseudoephedrine to methamphetamine. For this reason, red and white phosphorus were designated by the United States Drug Enforcement Administration as List I precursor chemicals under 21 CFR 1310.02 effective on November 17, 2001. In the United States, handlers of red or white phosphorus are subject to stringent regulatory controls.

Grow light (enhancing plant growth)

From Wikipedia, the free encyclopedia

A grow light or plant light is an artificial light source, generally an electric light, designed to stimulate plant growth by emitting a light appropriate for photosynthesis. Grow lights are used in applications where there is either no naturally occurring light, or where supplemental light is required. For example, in the winter months when the available hours of daylight may be insufficient for the desired plant growth, lights are used to extend the time the plants receive light. If plants do not receive enough light, they will grow long and spindly.

Grow lights either attempt to provide a light spectrum similar to that of the sun, or to provide a spectrum that is more tailored to the needs of the plants being cultivated. Outdoor conditions are mimicked with varying colour, temperatures and spectral outputs from the grow light, as well as varying the lumen output (intensity) of the lamps. Depending on the type of plant being cultivated, the stage of cultivation (e.g. the germination/vegetative phase or the flowering/fruiting phase), and the photoperiod required by the plants, specific ranges of spectrum, luminous efficacy and colour temperature are desirable for use with specific plants and time periods.

Russian botanist Andrei Famintsyn was the first to use artificial light for plant growing and research (1868).

Typical usage

Grow lights are used for horticulture, indoor gardening, plant propagation and food production, including indoor hydroponics and aquatic plants. Although most grow lights are used on an industrial level, they can also be used in households.

According to the inverse-square law, the intensity of light radiating from a point source (in this case a bulb) that reaches a surface is inversely proportional to the square of the surface's distance from the source (if an object is twice as far away, it receives only a quarter the light) which is a serious hurdle for indoor growers, and many techniques are employed to use light as efficiently as possible. Reflectors are thus often used in the lights to maximize light efficiency. Plants or lights are moved as close together as possible so that they receive equal lighting and that all light coming from the lights falls on the plants rather than on the surrounding area. 

Example of an HPS grow light set up in a grow tent. The setup includes a carbon filter to remove odors, and ducting to exhaust hot air using a powerful exhaust fan.

 

A range of bulb types can be used as grow lights, such as incandescents, fluorescent lights, high-intensity discharge lamps (HID), and light-emitting diodes (LED). Today, the most widely used lights for professional use are HIDs and fluorescents. Indoor flower and vegetable growers typically use high-pressure sodium (HPS/SON) and metal halide (MH) HID lights, but fluorescents and LEDs are replacing metal halides due to their efficiency and economy.

Metal halide lights are regularly used for the vegetative phase of plant growth, as they emit larger amounts of blue and ultraviolet radiation. With the introduction of ceramic metal halide lighting and full-spectrum metal halide lighting, they are increasingly being utilized as an exclusive source of light for both vegetative and reproductive growth stages. Blue spectrum light may trigger a greater vegetative response in plants.

High-pressure sodium lights are also used as a single source of light throughout the vegetative and reproductive stages. As well, they may be used as an amendment to full-spectrum lighting during the reproductive stage. Red spectrum light may trigger a greater flowering response in plants. If high-pressure sodium lights are used for the vegetative phase, plants grow slightly more quickly, but will have longer internodes, and may be longer overall.

In recent years LED technology has been introduced into the grow light market. By designing an indoor grow light using diodes, specific wavelengths of light can be produced. NASA has tested LED grow lights for their high efficiency in growing food in space for extraterrestrial colonization. Findings showed that plants are affected by light in the red, green and blue parts of the visible light spectrum.

Common types

High Intensity Discharge (HID) lights

While fluorescent lighting used to be the most common type of indoor grow light, HID lights are now the most popular. High intensity discharge lamps have a high lumen-per-watt efficiency. There are several different types of HID lights including mercury vapor, metal halide, high pressure sodium and conversion bulbs. Metal halide and HPS lamps produce a color spectrum that is somewhat comparable to the sun and can be used to grow plants. Mercury vapor lamps were the first type of HIDs and were widely used for street lighting, but when it comes to indoor gardening they produce a relatively poor spectrum for plant growth so they have been mostly replaced by other types of HIDs for growing plants.

All HID grow lights require a ballast to operate, and each ballast has a particular wattage. Popular HID wattages include 150W, 250W, 400W, 600W and 1000W. Of all the sizes, 600W HID lights are the most electrically efficient as far as light produced, followed by 1000W. A 600W HPS produces 7% more light (watt-for-watt) than a 1000W HPS.

Although all HID lamps work on the same principle, the different types of bulbs have different starting and voltage requirements, as well as different operating characteristics and physical shape. Because of this a bulb won't work properly unless it's using a matching ballast, even if the bulb will physically screw in. In addition to producing lower levels of light, mismatched bulbs and ballasts will stop working early, or may even burn out immediately.

Metal Halide (MH)

400W Metal halide bulb compared to smaller incandescent bulb

 

Metal halide bulbs are a type of HID light that emit light in the blue and violet parts of the light spectrum, which is similar to the light that is available outdoors during spring. Because their light mimics the color spectrum of the sun, some growers find that plants look more pleasing under a metal halide than other types of HID lights such as the HPS which distort the color of plants. Therefore, it's more common for a metal halide to be used when the plants are on display in the home (for example with ornamental plants) and natural color is preferred. Metal halide bulbs need to be replaced about once a year, compared to HPS lights which last twice as long.

Metal halide lamps are widely used in the horticultural industry and are well-suited to supporting plants in earlier developmental stages by promoting stronger roots, better resistance against disease and more compact growth. The blue spectrum of light encourages compact, leafy growth and may be better suited to growing vegetative plants with lots of foliage.

A metal halide bulb produces 60-125 lumens/watt, depending on the wattage of the bulb.

They are now being made for digital ballasts in a pulse start version, which have higher electrical efficiency (up to 110 lumens per watt) and faster warmup. One common example of a pulse start metal halide is the ceramic metal halide (CMH). Pulse start metal halide bulbs can come in any desired spectrum from cool white (7000 K) to warm white (3000 K) and even ultraviolet-heavy (10,000 K).

Ceramic Metal Halide (CMH, CDM)

Ceramic metal halide (CMH) lamps are a relatively new type of HID lighting, and the technology is referred to by a few names when it comes to grow lights, including ceramic discharge metal halide (CDM), ceramic arc metal halide. 

Ceramic metal halide lights are started with a pulse-starter, just like other "pulse-start" metal halides. The discharge of a ceramic metal halide bulb is contained in a type of ceramic material known as polycrystalline alumina (PCA), which is similar to the material used for an HPS. PCA reduces sodium loss, which in turn reduces color shift and variation compared to standard MH bulbs. Horticultural CDM offerings from companies such as Philips have proven to be effective sources of growth light for medium-wattage applications.

Combination MH and HPS ("Dual arc")

Combination HPS/MH lights combine a metal halide and a high-pressure sodium in the same bulb, providing both red and blue spectra in a single HID lamp. The combination of blue metal halide light and red high-pressure sodium light is an attempt to provide a very wide spectrum within a single lamp. This allows for a single bulb solution throughout the entire life cycle of the plant, from vegetative growth through flowering. There are potential trade-offs for the convenience of a single bulb in terms of yield. There are however some qualitative benefits that come for the wider light spectrum.

High-Pressure Sodium (HPS)

An HPS (High Pressure Sodium) grow light bulb in an air-cooled reflector with hammer finish. The yellowish light is the signature color produced by an HPS.

 

High-pressure sodium lights are a more efficient type of HID lighting than metal halides. HPS bulbs emit light in the yellow/red visible light as well as small portions of all other visible light. Since HPS grow lights deliver more energy in the red part of the light spectrum, they may promote blooming and fruiting. They are used as a supplement to natural daylight in greenhouse lighting and full-spectrum lighting(metal halide) or, as a standalone source of light for indoors/grow chambers. 

HPS grow lights are sold in the following sizes: 150W, 250W, 400W, 600W and 1000W. Of all the sizes, 600W HID lights are the most electrically efficient as far as light produced, followed by 1000W. A 600W HPS produces 7% more light (watt-for-watt) than a 1000W HPS.

A 600W High Pressure Sodium bulb

 

An HPS bulb produces 60-140 lumens/watt, depending on the wattage of the bulb.

Plants grown under HPS lights tend to elongate from the lack of blue/ultraviolet radiation. Modern horticultural HPS lamps have a much better adjusted spectrum for plant growth. The majority of HPS lamps while providing good growth, offer poor color rendering index (CRI) rendering. As a result, the yellowish light of an HPS can make monitoring plant health indoors more difficult. CRI isn't an issue when HPS lamps are used as supplemental lighting in greenhouses which make use of natural daylight (which offsets the yellow light of the HPS). 

High-pressure sodium lights have a long usable bulb life, and six times more light output per watt of energy consumed than a standard incandescent grow light. Due to their high efficiency and the fact that plants grown in greenhouses get all the blue light they need naturally, these lights are the preferred supplemental greenhouse lights. But, in the higher latitudes, there are periods of the year where sunlight is scarce, and additional sources of light are indicated for proper growth. HPS lights may cause distinctive infrared and optical signatures, which can attract insects or other species of pests; these may in turn threaten the plants being grown. High-pressure sodium lights emit a lot of heat, which can cause leggier growth, although this can be controlled by using special air-cooled bulb reflectors or enclosures.

Conversion bulbs

Conversion bulbs are manufactured so they work with either a MH or HPS ballast. A grower can run an HPS conversion bulb on a MH ballast, or a MH conversion bulb on a HPS ballast. The difference between the ballasts is an HPS ballast has an igniter which ignites the sodium in an HPS bulb, while a MH ballast does not. Because of this, all electrical ballasts can fire MH bulbs, but only a Switchable or HPS ballast can fire an HPS bulb without a conversion bulb. Usually a metal halide conversion bulb will be used in an HPS ballast since the MH conversion bulbs are more common.

Switchable ballasts

A switchable ballast is an HID ballast can be used with either a metal halide or an HPS bulb of equivalent wattage. So a 600W Switchable ballast would work with either a 600W MH or HPS. Growers use these fixtures for propagating and vegetatively growing plants under the metal halide, then switching to a high-pressure sodium bulb for the fruiting or flowering stage of plant growth. To change between the lights, only the bulb needs changing and a switch needs to be set to the appropriate setting.

LEDs (Light Emitting Diodes)

Two plants growing under an LED grow light

 

LED grow lights are composed of light-emitting diodes, usually in a casing with a heat sink and built-in fans. LED grow lights do not usually require a separate ballast and can be plugged directly into a standard electrical socket. 

LED grow lights vary in color depending on the intended use. It is known from the study of photomorphogenesis that green, red, far-red and blue light spectra have an effect on root formation, plant growth, and flowering, but there are not enough scientific studies or field-tested trials using LED grow lights to recommended specific color ratios for optimal plant growth under LED grow lights. It has been shown that many plants will grow normally if given both red and blue light. However, many studies indicate that red and blue light only provides the most cost efficient method of growth, plant growth is still better under light supplemented with green.

White LED grow lights provide a full spectrum of light designed to mimic natural light, providing plants a balanced spectrum of red, blue and green. The spectrum used varies, however, white LED grow lights are designed to emit similar amounts of red and blue light with the added green light to appear white. White LED grow lights are often used for supplemental lighting in home and office spaces. 

A large number of plant species have been assessed in greenhouse trials to make sure plants have higher quality in biomass and biochemical ingredients even higher or comparable with field conditions. Plant performance of mint, basil, lentil, lettuce, cabbage, parsley, carrot were measured by assessing health and vigor of plants and success in promoting growth. Promoting in profuse flowering of select ornamentals including primula, marigold, stock were also noticed.

In tests conducted by Philips Lighting on LED grow lights to find an optimal light recipe for growing various vegetables in greenhouses, they found that the following aspects of light affects both plant growth (photosynthesis) and plant development (morphology): light intensity, total light over time, light at which moment of the day, light/dark period per day, light quality (spectrum), light direction and light distribution over the plants. However it's noted that in tests between tomatoes, mini cucumbers and bell peppers, the optimal light recipe was not the same for all plants, and varied depending on both the crop and the region, so currently they must optimize LED lighting in greenhouses based on trial and error. They've shown that LED light affects disease resistance, taste and nutritional levels, but as of 2014 they haven't found a practical way to use that information.

Ficus plant grown under a white LED grow light.

 

The diodes used in initial LED grow light designs were usually 1/3 watt to 1 watt in power. However, higher wattage diodes such as 3 watt and 5 watt diodes are now commonly used in LED grow lights. for highly compacted areas, COB chips between 10 watts and 100 watts can be used. Because of heat dissipation, these chips are often less efficient. 

LED grow lights should be kept at least 12 inches (30 cm) away from plants to prevent leaf burn.

Historically, LED lighting was very expensive, but costs have greatly reduced over time, and their longevity has made them more popular. LED grow lights are often priced higher, watt-for-watt, than other LED lighting, due to design features that help them to be more energy efficient and last longer. In particular, because LED grow lights are relatively high power, LED grow lights are often equipped with cooling systems, as low temperature improves both the brightness and longevity. LEDs usually last for 50,000 - 90,000 hours until LM-70 is reached.

Fluorescent

Fluorescent grow light

 

Fluorescent lights come in many form factors, including long, thin bulbs as well as smaller spiral shaped bulbs (compact fluorescent lights). Fluorescent lights are available in color temperatures ranging from 2700 K to 10,000 K. The luminous efficacy ranges from 30 lm/W to 90 lm/W. The two main types of fluorescent lights used for growing plants are the tube-style lights and compact fluorescent lights.

Tube-style fluorescent lights

Fluorescent grow lights are not as intense as HID lights and are usually used for growing vegetables and herbs indoors, or for starting seedlings to get a jump start on spring plantings. A ballast is needed to run these types of fluorescent lights.

Standard fluorescent lighting comes in multiple form factors, including the T5, T8 and T12. The brightest version is the T5. The T8 and T12 are less powerful and are more suited to plants with lower light needs. High-output fluorescent lights produce twice as much light as standard fluorescent lights. A high-output fluorescent fixture has a very thin profile, making it useful in vertically limited areas. 

Fluorescents have an average usable life span of up to 20,000 hours. A fluorescent grow light produces 33-100 lumens/watt, depending on the form factor and wattage.

Compact Fluorescent Lights (CFLs)

Dual spectrum compact fluorescent grow light. Actual length is about 40 cm (16 in)

 

Standard Compact Fluorescent Light

 

Compact Fluorescent lights (CFLs) are smaller versions of fluorescent lights that were originally designed as pre-heat lamps, but are now available in rapid-start form. CFLs have largely replaced incandescent light bulbs in households because they last longer and are much more electrically efficient. In some cases, CFLs are also used as grow lights. Like standard fluorescent lights, they are useful for propagation and situations where relatively low light levels are needed. 

While standard CFLs in small sizes can be used to grow plants, there are also now CFL lamps made specifically for growing plants. Often these larger compact fluorescent bulbs are sold with specially designed reflectors that direct light to plants, much like HID lights. Common CFL grow lamp sizes include 125W, 200W, 250W and 300W.

Unlike HID lights, CFLs fit in a standard mogul light socket and don't need a separate ballast.

Compact fluorescent bulbs are available in warm/red (2700 K), full spectrum or daylight (5000 K) and cool/blue (6500 K) versions. Warm red spectrum is recommended for flowering, and cool blue spectrum is recommended for vegetative growth.

Usable life span for compact fluorescent grow lights is about 10,000 hours. A CFL produces 44-80 lumens/watt, depending on the wattage of the bulb.

Examples of lumens and lumens/watt for different size CFLs:

CFL Wattage	Initial Lumens	Lumens/watt
23W	1,600	70
42W	2,800	67
85W	4,250	50
125W	7,000	56
200W	10,000	50

 Cold Cathode Fluorescent Light (CCFL)

A cold cathode is a cathode that is not electrically heated by a filament. A cathode may be considered "cold" if it emits more electrons than can be supplied by thermionic emissionalone. It is used in gas-discharge lamps, such as neon lamps, discharge tubes, and some types of vacuum tube. The other type of cathode is a hot cathode, which is heated by electric current passing through a filament. A cold cathode does not necessarily operate at a low temperature: it is often heated to its operating temperature by other methods, such as the current passing from the cathode into the gas.

Color spectrum

The color temperatures of different grow lights

Different grow lights produce different spectrums of light. Plant growth patterns can respond to the color spectrum of light, a process completely separate from photosynthesis known as photomorphogenesis.

Natural daylight has a high color temperature (approximately 5000-5800 K). Visible light color varies according to the weather and the angle of the Sun, and specific quantities of light (measured in lumens) stimulate photosynthesis. Distance from the sun has little effect on seasonal changes in the quality and quantity of light and the resulting plant behavior during those seasons. The axial tilt of the Earth is not perpendicular to the plane of its orbit around the sun. During half of the year the north pole is tilted towards sun so the northern hemisphere gets nearly direct sunlight and the southern hemisphere gets oblique sunlight that must travel through more atmosphere before it reaches the Earth's surface. In the other half of the year, this is reversed. The color spectrum of visible light that the sun emits does not change, only the quantity (more during the summer and less in winter) and quality of overall light reaching the Earth's surface. Some supplemental LED grow lights in vertical greenhouses produce a combination of only red and blue wavelengths. The color rendering index facilitates comparison of how closely the light matches the natural color of regular sunlight. 

The ability of a plant to absorb light varies with species and environment, however, the general measurement for the light quality as it affects plants is the PAR value, or Photosynthetically Active Radiation.

There have been several experiments using LEDs to grow plants, and it has been shown that plants need both red and blue light for healthy growth. From experiments it has been consistently found that the plants that are growing only under LEDs red (660 nm, long waves) spectrum growing poorly with leaf deformities, though adding a small amount of blue allows most plants to grow normally.

Several reports suggest that a minimum blue light requirement of 15-30 µmol·m⁻²·s⁻¹ is necessary for normal development in several plant species.

LED panel light source used in an experiment on potato plant growth by NASA

 

Many studies indicate that even with blue light added to red LEDs, plant growth is still better under white light, or light supplemented with green. Neil C Yorio demonstrated that by adding 10% blue light (400 to 500 nm) to the red light (660 nm) in LEDs, certain plants like lettuce and wheat grow normally, producing the same dry weight as control plants grown under full spectrum light. However, other plants like radish and spinach grow poorly, and although they did better under 10% blue light than red-only light, they still produced significantly lower dry weights compared to control plants under a full spectrum light. Yorio speculates there may be additional spectra of light that some plants need for optimal growth.

Greg D. Goins examined the growth and seed yield of Arabidopsis plants grown from seed to seed under red LED lights with 0%, 1%, or 10% blue spectrum light. Arabidopsis plants grown under only red LEDS alone produced seeds, but had unhealthy leaves, and plants took twice as long to start flowering compared to the other plants in the experiment that had access to blue light. Plants grown with 10% blue light produced half the seeds of those grown under full spectrum, and those with 0% or 1% blue light produced one-tenth the seeds of the full spectrum plants. The seeds all germinated at a high rate under all light types tested.

Hyeon-Hye Kim demonstrated that the addition of 24% green light (500-600 nm) to red and blue LEDs enhanced the growth of lettuce plants. These RGB treated plants not only produced higher dry and wet weight and greater leaf area than plants grown under just red and blue LEDs, they also produced more than control plants grown under cool white fluorescent lamps, which are the typical standard for full spectrum light in plant research. She reported that the addition of green light also makes it easier to see if the plant is healthy since leaves appear green and normal. However, giving nearly all green light (86%) to lettuce produced lower yields than all the other groups.

The National Aeronautics and Space Administration’s (NASA) Biological Sciences research group has concluded that light sources consisting of more than 50% green cause reductions in plant growth, whereas combinations including up to 24% green enhance growth for some species. Green light has been shown to affect plant processes via both cryptochrome-dependent and cryptochrome-independent means. Generally, the effects of green light are the opposite of those directed by red and blue wavebands, and it's speculated that green light works in orchestration with red and blue.

Light requirements of plants

Absorbance spectra of free chlorophyll a (blue) and b (red) in a solvent. The action spectra of chlorophyll molecules are slightly modified in vivo depending on specific pigment-protein interactions.

 

A plant's specific needs determine which lighting is most appropriate for optimum growth. If a plant does not get enough light, it will not grow, regardless of other conditions. Most plants use chlorophyll which mostly reflects green light, but absorbs red and blue light well. Vegetables grow best in strong sunlight, and to flourish indoors they need sufficient light levels, whereas foliage plants (e.g. Philodendron) grow in full shade and can grow normally with much lower light levels. 

Grow lights usage is dependent on the plant's phase of growth. Generally speaking, during the seedling/clone phase, plants should receive 16+ hours on, 8- hours off. The vegetative phase typically requires 18 hours on, and 6 hours off. During the final, flower stage of growth, keeping grow lights on for 12 hours on and 12 hours off is recommended.

Photoperiodism

In addition, many plants also require both dark and light periods, an effect known as photoperiodism, to trigger flowering. Therefore, lights may be turned on or off at set times. The optimum photo/dark period ratio depends on the species and variety of plant, as some prefer long days and short nights and others prefer the opposite or intermediate "day lengths". 

Much emphasis is placed on photoperiod when discussing plant development. However, it is the number of hours of darkness that affects a plant’s response to day length. In general, a “short-day” is one in which the photoperiod is no more than 12 hours. A “long-day” is one in which the photoperiod is no less than 14 hours. Short-day plants are those that flower when the day length is less than a critical duration. Long-day plants are those that only flower when the photoperiod is greater than a critical duration. Day-neutral plants are those that flower regardless of photoperiod.

Plants that flower in response to photoperiod may have a facultative or obligate response. A facultative response means that a plant will eventually flower regardless of photoperiod, but will flower faster if grown under a particular photoperiod. An obligate response means that the plant will only flower if grown under a certain photoperiod.

Photosynthetically Active Radiation (PAR)

Weighting factor for photosynthesis. The photon-weighted curve is for converting PPFD to YPF; the energy-weighted curve is for weighting PAR expressed in watts or joules.

 

Lux and lumens are commonly used to measure light levels, but they are photometric units which measure the intensity of light as perceived by the human eye.

The spectral levels of light that can be used by plants for photosynthesis is similar to, but not the same as what's measured by lumens. Therefore, when it comes to measuring the amount of light available to plants for photosynthesis, biologists often measure the amount of photosynthetically active radiation (PAR) received by a plant. PAR designates the spectral range of solar radiation from 400 to 700 nanometers, which generally corresponds to the spectral range that photosynthetic organisms are able to use in the process of photosynthesis. 

The irradiance of PAR can be expressed in units of energy flux (W/m²), which is relevant in energy-balance considerations for photosynthetic organisms. However, photosynthesis is a quantum process and the chemical reactions of photosynthesis are more dependent on the number of photons than the amount of energy contained in the photons. Therefore, plant biologists often quantify PAR using the number of photons in the 400-700 nm range received by a surface for a specified amount of time, or the Photosynthetic Photon Flux Density (PPFD). This is normally measured using mol m⁻²s⁻¹. 

According to one manufacturer of grow lights, plants require at least light levels between 100 and 800 μmol m⁻²s⁻¹. For daylight-spectrum (5800 K) lamps, this would be equivalent to 5800 to 46,000 lm/m².

Chlorophyll

From Wikipedia, the free encyclopedia

Chlorophyll at different scales

Chlorophyll is responsible for the green color of many plants and algae.

 

Seen through a microscope, chlorophyll is concentrated within organisms in structures called chloroplasts.

 

Plants are perceived as green because chlorophyll absorbs mainly the blue and red wavelenght and reflects the green.

 

There are several types of chlorophyll, but all share the chlorin magnesium ligand which forms the right side of this diagram.

Chlorophyll is any of several related green pigments found in cyanobacteria and the chloroplasts of algae and plants. Its name is derived from the Greek words χλωρός, chloros ("green") and φύλλον, phyllon ("leaf"). Chlorophyll is essential in photosynthesis, allowing plants to absorb energy from light.

Chlorophylls absorb light most strongly in the blue portion of the electromagnetic spectrum as well as the red portion. Conversely, it is a poor absorber of green and near-green portions of the spectrum, which it reflects, producing the green color of chlorophyll-containing tissues. Two types of chlorophyll exist in the photosystems of green plants: chlorophyll a and b.

History

Chlorophyll was first isolated and named by Joseph Bienaimé Caventou and Pierre Joseph Pelletier in 1817. The presence of magnesium in chlorophyll was discovered in 1906, and was the first time that magnesium had been detected in living tissue.

After initial work done by German chemist Richard Willstätter spanning from 1905 to 1915, the general structure of chlorophyll a was elucidated by Hans Fischer in 1940. By 1960, when most of the stereochemistry of chlorophyll a was known, Robert Burns Woodward published a total synthesis of the molecule. In 1967, the last remaining stereochemical elucidation was completed by Ian Fleming, and in 1990 Woodward and co-authors published an updated synthesis. Chlorophyll f was announced to be present in cyanobacteria and other oxygenic microorganisms that form stromatolites in 2010; a molecular formula of C₅₅H₇₀O₆N₄Mg and a structure of (2-formyl)-chlorophyll a were deduced based on NMR, optical and mass spectra.

Photosynthesis

Absorbance spectra of free chlorophyll a (blue) and b (red) in a solvent. The spectra of chlorophyll molecules are slightly modified in vivo depending on specific pigment-protein interactions.

 

Chlorophyll is vital for photosynthesis, which allows plants to absorb energy from light.

Chlorophyll molecules are arranged in and around photosystems that are embedded in the thylakoid membranes of chloroplasts. In these complexes, chlorophyll serves three functions. The function of the vast majority of chlorophyll (up to several hundred molecules per photosystem) is to absorb light. Having done so, these same centers execute their second function: the transfer of that light energy by resonance energy transfer to a specific chlorophyll pair in the reaction center of the photosystems. This pair effects the final function of chlorophylls, charge separation, leading to biosynthesis. The two currently accepted photosystem units are photosystem II and photosystem I, which have their own distinct reaction centers, named P680 and P700, respectively. These centres are named after the wavelength (in nanometers) of their red-peak absorption maximum. The identity, function and spectral properties of the types of chlorophyll in each photosystem are distinct and determined by each other and the protein structure surrounding them. Once extracted from the protein into a solvent (such as acetone or methanol), these chlorophyll pigments can be separated into chlorophyll a and chlorophyll b. 

The function of the reaction center of chlorophyll is to absorb light energy and transfer it to other parts of the photosystem. The absorbed energy of the photon is transferred to an electron in a process called charge separation. The removal of the electron from the chlorophyll is an oxidation reaction. The chlorophyll donates the high energy electron to a series of molecular intermediates called an electron transport chain. The charged reaction center of chlorophyll (P680⁺) is then reduced back to its ground state by accepting an electron stripped from water. The electron that reduces P680⁺ ultimately comes from the oxidation of water into O₂ and H⁺ through several intermediates. This reaction is how photosynthetic organisms such as plants produce O₂ gas, and is the source for practically all the O₂ in Earth's atmosphere. Photosystem I typically works in series with Photosystem II; thus the P700⁺ of Photosystem I is usually reduced as it accepts the electron, via many intermediates in the thylakoid membrane, by electrons coming, ultimately, from Photosystem II. Electron transfer reactions in the thylakoid membranes are complex, however, and the source of electrons used to reduce P700⁺ can vary. 

The electron flow produced by the reaction center chlorophyll pigments is used to pump H⁺ ions across the thylakoid membrane, setting up a chemiosmotic potential used mainly in the production of ATP (stored chemical energy) or to reduce NADP⁺ to NADPH. NADPH is a universal agent used to reduce CO₂ into sugars as well as other biosynthetic reactions.

Reaction center chlorophyll–protein complexes are capable of directly absorbing light and performing charge separation events without the assistance of other chlorophyll pigments, but the probability of that happening under a given light intensity is small. Thus, the other chlorophylls in the photosystem and antenna pigment proteins all cooperatively absorb and funnel light energy to the reaction center. Besides chlorophyll a, there are other pigments, called accessory pigments, which occur in these pigment–protein antenna complexes.

Chemical structure

Space-filling model of the chlorophyll a molecule

 

Chlorophylls are numerous in types, but all are defined by the presence of a fifth ring beyond the four pyrrole-like rings. Most chlorophylls are classified as chlorins, which are reduced relatives to porphyrins (found in hemoglobin). They share a common biosynthetic pathway as porphyrins, including the precursor uroporphyrinogen III. Unlike hemes, which feature iron at the center of the tetrapyrrole ring, chlorophylls bind magnesium. For the structures depicted in this article, some of the ligands attached to the Mg²⁺ center are omitted for clarity. The chlorin ring can have various side chains, usually including a long phytol chain. The most widely distributed form in terrestrial plants is chlorophyll a. 

When leaves degreen in the process of plant senescence, chlorophyll is converted to a group of colorless tetrapyrroles known as nonfluorescent chlorophyll catabolites (NCC's) with the general structure: 

These compounds have also been identified in several ripening fruits.

Measurement of chlorophyll content

Pure chlorophyll has a strong green color.

Measurement of the absorption of light is complicated by the solvent used to extract the chlorophyll from plant material, which affects the values obtained,

• In diethyl ether, chlorophyll a has approximate absorbance maxima of 430 nm and 662 nm, while chlorophyll b has approximate maxima of 453 nm and 642 nm.
• The absorption peaks of chlorophyll a are at 465 nm and 665 nm. Chlorophyll a fluoresces at 673 nm (maximum) and 726 nm. The peak molar absorption coefficient of chlorophyll a exceeds 10⁵ M⁻¹ cm⁻¹, which is among the highest for small-molecule organic compounds.
• In 90% acetone-water, the peak absorption wavelengths of chlorophyll a are 430 nm and 664 nm; peaks for chlorophyll b are 460 nm and 647 nm; peaks for chlorophyll c1 are 442 nm and 630 nm; peaks for chlorophyll c2 are 444 nm and 630 nm; peaks for chlorophyll d are 401 nm, 455 nm and 696 nm.

By measuring the absorption of light in the red and far red regions, it is possible to estimate the concentration of chlorophyll within a leaf.

Ratio fluorescence emission can be used to measure chlorophyll content. By exciting chlorophyll a fluorescence at a lower wavelength, the ratio of chlorophyll fluorescence emission at 705±10 nm and 735±10 nm can provide a linear relationship of chlorophyll content when compared to chemical testing. The ratio F₇₃₅/F₇₀₀ provided a correlation value of r² 0.96 compared to chemical testing in the range from 41 mg m⁻² up to 675 mg m⁻². Gitelson also developed a formula for direct readout of chlorophyll content in mg m⁻². The formula provided a reliable method of measuring chlorophyll content from 41 mg m⁻² up to 675 mg m⁻² with a correlation r² value of 0.95.

Biosynthesis

In plants, chlorophyll may be synthesized from succinyl-CoA and glycine, although the immediate precursor to chlorophyll a and b is protochlorophyllide. In Angiosperm plants, the last step, the conversion of protochlorophyllide to chlorophyll, is light-dependent and such plants are pale (etiolated) if grown in darkness. Non-vascular plants and green algae have an additional light-independent enzyme and grow green even in darkness. 

Chlorophyll itself is bound to proteins and can transfer the absorbed energy in the required direction. Protochlorophyllide occurs mostly in the free form and, under light conditions, acts as a photosensitizer, forming highly toxic free radicals. Hence, plants need an efficient mechanism of regulating the amount of chlorophyll precursor. In angiosperms, this is done at the step of aminolevulinic acid (ALA), one of the intermediate compounds in the biosynthesis pathway. Plants that are fed by ALA accumulate high and toxic levels of protochlorophyllide; so do the mutants with the damaged regulatory system.

Chlorosis is a condition in which leaves produce insufficient chlorophyll, turning them yellow. Chlorosis can be caused by a nutrient deficiency of iron — called iron chlorosis — or by a shortage of magnesium or nitrogen. Soil pH sometimes plays a role in nutrient-caused chlorosis; many plants are adapted to grow in soils with specific pH levels and their ability to absorb nutrients from the soil can be dependent on this. Chlorosis can also be caused by pathogens including viruses, bacteria and fungal infections, or sap-sucking insects.

Complementary light absorbance of anthocyanins with chlorophylls

Superposition of spectra of chlorophyll a and b with oenin (malvidin 3O glucoside), a typical anthocyanidin, showing that, while chlorophylls absorb in the blue and yellow/red parts of the visible spectrum, oenin absorbs mainly in the green part of the spectrum, where chlorophylls don't absorb at all.

 

Anthocyanins are other plant pigments. The absorbance pattern responsible for the red color of anthocyanins may be complementary to that of green chlorophyll in photosynthetically active tissues such as young Quercus coccifera leaves. It may protect the leaves from attacks by plant eaters that may be attracted by green color.

Distribution

The chlorophyll maps show milligrams of chlorophyll per cubic meter of seawater each month. Places where chlorophyll amounts were very low, indicating very low numbers of phytoplankton, are blue. Places where chlorophyll concentrations were high, meaning many phytoplankton were growing, are yellow. The observations come from the Moderate Resolution Imaging Spectroradiometer (MODIS) on NASA's Aqua satellite. Land is dark gray, and places where MODIS could not collect data because of sea ice, polar darkness, or clouds are light gray.The highest chlorophyll concentrations, where tiny surface-dwelling ocean plants are thriving, are in cold polar waters or in places where ocean currents bring cold water to the surface, such as around the equator and along the shores of continents. It is not the cold water itself that stimulates the phytoplankton. Instead, the cool temperatures are often a sign that the water has welled up to the surface from deeper in the ocean, carrying nutrients that have built up over time. In polar waters, nutrients accumulate in surface waters during the dark winter months when plants cannot grow. When sunlight returns in the spring and summer, the plants flourish in high concentrations.

Culinary use

Chlorophyll is registered as a food additive (colorant), and its E number is E140. Chefs use chlorophyll to color a variety of foods and beverages green, such as pasta and spirits. Absinthe gains its green color from the chlorophyll introduced through the large variety of herbs included in its production. Chlorophyll is not soluble in water, and it is first mixed with a small quantity of vegetable oil to obtain the desired solution.