1. Phosphorus Atomic Number
2. Phosphorus Atom Valence Electrons

In biological molecules, phosphorus is usually found in organophosphates. Organophosphates are made up of a phosphorus atom bonded to four oxygens, with one of the oxygens also bonded to a carbon. In methyl phosphate, the phosphorus is sp3 hybridized and the O-P-O bond angle varies from 110° to 112 o. White phosphorus has also found a wide range of other uses. One of these was in phosphorus matches that were first sold in Stockton-on-Tees in the UK in 1827. This created a whole new industry of cheap lights - but at a terrible cost. Breathing in phosphorus vapour led to the industrial disease phossy jaw, which slowly ate away the jaw bone. P I Ground State 1s 2 2s 2 2p 6 3s 2 3p 3 4 S° 3 / 2 Ionization energy 84580.83 cm-1 (10.48669 eV) Ref. MZM85 P II Ground State 1s 2 2s 2 2p 6 3s 2 3p 2 3 P 0 Ionization energy 159451.5 cm-1 (19.7695 eV) Ref. What is Phosphorus Phosphorus is a chemical element with atomic number 15 which means there are 15 protons and 15 electrons in the atomic structure. The chemical symbol for Phosphorus is P. From its atomic number of 15, it is possible to predict that the phosphorus atom has. A) 5 neutrons, 5 protons, and 5 electrons B) 15 neutrons and 15 protons C) 8 electrons in its outermost electron shell.

Phosphorus is a chemical element in the periodic table that has the symbol P and atomic number 15. A multivalent, nonmetal of the nitrogen group, phosphorus is commonly found in inorganic phosphate rocks and in all living cells but is never naturally found alone. It is highly reactive, emits a faint glow upon uniting with oxygen (hence its name, Latin for 'morning star', from Greek words meaning 'light' and 'bring'), occurs in several forms, and is an essential element for living organisms. The most important use of phosphorus is in the production of fertilizers. It is also widely used in explosives, friction matches, fireworks, pesticides, toothpaste, and detergents.

Notable characteristics

Common phosphorus forms a waxy white solid that has a characteristic disagreeable smell but when it is pure it is colorless and transparent. This non metal is not soluble in water, but it is soluble in carbon disulfide. Pure phosphorus ignites spontaneously in air and burns to phosphorus pentoxide.

Forms

Phosphorus exists in four or more allotropic forms: white (or yellow), red, and black (or violet). The most common are red and white phosphorus, both of which are tetrahedral groups of four atoms. White phosphorus burns on contact with air and on exposure to heat or light it can transform into red phosphorus. It also exists in two modifications: alpha and beta which are separated by a transition temperature of -3.8 °C. Red phosphorus is comparatively stable and sublimes at a vapor pressure of 1 atm at 170 °C but burns from impact or frictional heating. A black phosphorus allotrope exists which has a structure similar to graphite - the atoms are arranged in hexagonal sheet layers and will conduct electricity.

Applications

Concentrated phosphoric acids, which can consist of 70% to 75% P₂O₅ are very important to agriculture and farm production in the form of fertilizers. Global demand for fertilizers has led to large increases in phosphate production in the second half of the 20th century. Other uses;

• Phosphates are utilized in the making of special glasses that are used for sodium lamps.
• Bone-ash, calcium phosphate, is used in the production of fine china and to make mono-calcium phosphate which is employed in baking powder.
• This element is also an important component in steel production, n themaking of phosphor bronze, and in many other related products.
• Trisodium phosphate is widely used in cleaning agents to soften water and for preventing pipe/boiler tube corrosion.
• White phosphorus is used in military applications as incendiary bombs, smoke pots, smoke bombs and tracer bullets.
• Miscellaneous uses; used in the making of safety matches, pyrotechnics, pesticides, toothpaste, detergents, etc.

Biological role

Phosphorus compounds perform vital functions in all known forms of life. Inorganic phosphorus plays a key role in biological molecules such as DNA and RNA where it forms part of those molecules' molecular backbones. Living cells also utilize inorganic phosphorus to store and transport cellular energy via adenosine triphosphate (ATP). Calcium phosphate salts are used by animals to stiffen bones and phosphorus is also an important element in cell protoplasm and nervous tissue.

History

Phosphorus (Greek. phosphoros, meaning 'light bearer' which was the ancient name for the planet Venus) was discovered by German alchemist Hennig Brand in 1669 through a preparation from urine. Working in Hamburg, Brand attempted to distill salts by evaporating urine, and in the process produced a white material that glowed in the dark and burned brilliantly. Since that time, phosphorescence has been used to describe substances that shine in the dark without burning.

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 giving off luminous steam). In addition, exposure to the vapors gave match workers a necrosis of the bones of the jaw, the infamous 'phossy-jaw.' When red phosphorus was discovered, with its far lower flammability and toxicity, it was adopted as a safer alternative for match manufacture.

Occurrence

Due to its reactivity to air and many other oxygen containing substances, phosphorus is not found free in nature but it is widely distributed in many different minerals. Phosphate rock, which is partially made of apatite (an impure tri-calcium phosphate mineral) is an important commercial source of this element. Large deposits of apatite are in Russia, Morocco, Florida, Idaho, Tennessee, Utah, and elsewhere.

The white allotrope can be produced using several different methods. In one process, tri-calcium phosphate, which is derived from phosphate rock, is heated in an electric or fuel-fired furnace in the presence of carbon and silica. Elemental phosphorus is then liberated as a vapor and can be collected under phosphoric acid.

Precautions

This is a particularly poisonous element with 50 mg being the average fatal dose. The allotrope white phosphorus should be kept under water at all times due to its hyper reactivity to air and it should only be manipulated with forceps since contact with skin can cause severe burns. Chronic white phosphorus poisoning of unprotected workers leads to necrosis of the jaw called 'phossy-jaw'. Phosphate esters are nerve poisons but inorganic phosphates are relatively nontoxic. Phosphate pollution occurs where fertilizers or detergents have leached into soils.

When the white form is exposed to sunlight or when it is heated in its own vapor to 250 °C, it is transmuted to the red form, which does not phosphoresce in air. The red allotrope does not spontaneously ignite in air and is not as dangerous as the white form. Nevertheless, it should be handled with care because it does revert to white phosphorus in some temperature ranges and it also emits highly toxic fumes that consist of phosphorus oxides when it is heated.

Isotopes

Some common isotopes of phosphorus include:

• ³²P (radioactive). Phosphorus-32 is a beta-emitter (1.71 MeV) with a half-life of 14.3 days. It is used routinely in life-science laboratories, primarily to produce radiolabeled DNA and RNA probes.

• ³³P (radioactive). Phosphorus-33 is 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.

Spelling

The only correct spelling of the element is phosphorus. There does exist a word phosphorous, but it is the adjectival form for the smaller valency: so, just as sulfur forms sulfurous and sulfuric compounds, so phosphorus forms phosphorous and phosphoric compounds.

Reference

• Los Alamos National Laboratory – Phosphorus(http://periodic.lanl.gov/elements/15.html)

External links

Phosphorus Atomic Number

• WebElements.com – Phosphorus(http://www.webelements.com/webelements/elements/text/P/index.html)
• EnvironmentalChemistry.com – Phosphorus(http://environmentalchemistry.com/yogi/periodic/P.html)

Phosphorus is a chemical element found on Earth in numerous compound forms, such as the phosphate ion (PO₄^3-), located in water, soil and sediments. The quantities of phosphorus in soil are generally small, and this often limits plant growth. That is why people often apply phosphate fertilisers on farmland. Animals absorb phosphates by eating plants or plant-eating animals.

The role of phosphorus in animals and plants

Phosphorus is an essential nutrient for animals and plants. It plays a critical role in cell development and is a key component of molecules that store energy, such as ATP (adenosine triphosphate), DNA and lipids (fats and oils). Insufficient phosphorus in the soil can result in a decreased crop yield.

The phosphorus cycle

Phosphorus moves in a cycle through rocks, water, soil and sediments and organisms.

Here are the key steps of the phosphorus cycle

• Over time, rain and weathering cause rocks to release phosphate ions and other minerals. This inorganic phosphate is then distributed in soils and water.
• Plants take up inorganic phosphate from the soil. The plants may then be consumed by animals. Once in the plant or animal, the phosphate is incorporated into organic molecules such as DNA. When the plant or animal dies, it decays, and the organic phosphate is returned to the soil.
• Within the soil, organic forms of phosphate can be made available to plants by bacteria that break down organic matter to inorganic forms of phosphorus. This process is known as mineralisation.
• Phosphorus in soil can end up in waterways and eventually oceans. Once there, it can be incorporated into sediments over time.

Most phosphorus is unavailable to plants

Since most of our phosphorus is locked up in sediments and rocks, it’s not available for plants to use. A lot of the phosphorus in soils is also not available to plants.

The availability of phosphorus in soil to plants depends of several reversible pathways:

• Bacteria: Bacteria convert plant-available phosphate into organic forms that are then not available to plants. Although other bacteria make phosphate available by mineralisation, the contribution of this is small.
• Adsorption: Inorganic (and available) phosphorus can be chemically bound (adsorbed) to soil particles, making it unavailable to plants. Desorption is the release of adsorbed phosphorus from its bound state into soil solution.
• pH: Inorganic phosphorus compounds need to be soluble to be taken up by plants. This depends on the acidity (pH) of the soil. If soils are less than pH 4 or greater than pH 8, the phosphorus starts to become tied up with other compounds, making it less available to plants.

Many plant crops need more phosphorus than is dissolved in the soil to grow optimally. In addition, crops are usually harvested and removed – leaving no decaying vegetation to replace phosphorus. Therefore, farmers replenish the phosphorus ‘pool’ by adding fertilisers or effluent to replace the phosphorus taken up by plants.

Phosphate fertilisers replenish soil phosphorus

Many farmers replenish phosphorus through the use of phosphate fertilisers. The phosphorus is obtained by mining deposits of rock phosphate. Locally produced sulfuric acid is used to convert the insoluble rock phosphate into a more soluble and usable form – a fertiliser product called superphosphate.

In New Zealand, superphosphate is made using rock imported mainly from Morocco.

Adjusting the pH of the soil for efficient plant uptake of phosphate should be done prior to fertilisation. For example, adding lime reduces soil acidity, which provides an environment where phosphate becomes more available to plants.

Water pollution by fertilisers

When fields are overfertilised (through commercial fertilisers or manure), phosphate not utilised by plants can be lost from the soil through leaching and water run-off. This phosphate ends up in waterways, lakes and estuaries. Excess phosphate causes excessive growth of plants in waterways, lakes and estuaries leading to eutrophication.

Steps are being taken in agriculture to reduce phosphate losses in order to maximise the efficiency of fertiliser and effluent applications.

Phosphorus Atom Valence Electrons

Nature of science

Scientists make observations and develop their explanations using inference, imagination and creativity. Often they use models to help other scientists understand their theories. The phosphorus cycle diagram is an example of an explanatory model. Diagrams demonstrate the creativity required by scientists to use their observations to develop models and to communicate their explanations to others.