Phosphorus Atom

1. Phosphorus Atomic Number
2. Phosphorus Atoms
3. Phosphorus Atom Diagram
4. Phosphorus Atom Model
5. Phosphorus Atomic Structure
6. Phosphorus Atomic Size

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.

Phosphorus atoms have fifteen protons and sixteen neutrons, so they're just a little lighter than sulphur atoms, in the middle of the range for atoms. Stars make phosphorus when two oxygen atoms get very hot and squash together into one phosphorus atom. 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. Black phosphorus has an orthorhombic pleated honeycomb structure and is the least reactive allotrope, a result of its lattice of interlinked six-membered rings where each atom is bonded to three other atoms. Black and red phosphorus can also take a cubic crystal lattice structure.

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

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

Objective

After completing this section, you should be able to apply the concept of hybridization of atoms such as N, O, P and S to explain the structures of simple species containing these atoms.

Key Terms

Make certain that you can define, and use in context, the key term below.

• lone pair electrons

Study Notes

Nitrogen is frequently found in organic compounds. As with carbon atoms, nitrogen atoms can be sp³-, sp²- or sp‑hybridized.

Note that, in this course, the term 'lone pair' is used to describe an unshared pair of electrons.

The valence-bond concept of orbital hybridization can be extrapolated to other atoms including nitrogen, oxygen, phosphorus, and sulfur. In other compounds, covalent bonds that are formed can be described using hybrid orbitals.

Nitrogen

Bonding in NH₃

The nitrogen in NH₃ has five valence electrons. After hybridization these five electrons are placed in the four equivalent sp³ hybrid orbitals. The electron configuration of nitrogen now has one sp³ hybrid orbital completely filled with two electrons and three sp³ hybrid orbitals with one unpaired electron each. The two electrons in the filled sp³ hybrid orbital are considered non-bonding because they are already paired. These electrons will be represented as a lone pair on the structure of NH₃. The three unpaired electrons in the hybrid orbitals are considered bonding and will overlap with the s orbitals in hydrogen to form N-H sigma bonds. Note! This bonding configuration was predicted by the Lewis structure of NH₃.

The four sp³ hybrid orbitals of nitrogen orientate themselves to form a tetrahedral geometry. The three N-H sigma bonds of NH₃ are formed by sp³(N)-1s(H) orbital overlap. The fourth sp³ hybrid orbital contains the two electrons of the lone pair and is not directly involved in bonding.

Phosphorus Atomic Number

Methyl amine

The nitrogen is sp³hybridized which means that it has four sp³ hybrid orbitals. Two of the sp³hybridized orbitals overlap with s orbitals from hydrogens to form the two N-H sigma bonds. One of the sp³hybridized orbitals overlap with an sp³ hybridized orbital from carbon to form the C-N sigma bond. The lone pair electrons on the nitrogen are contained in the last sp³ hybridized orbital. Due to the sp³hybridization the nitrogen has a tetrahedral geometry. However, the H-N-H and H-N-C bonds angles are less than the typical 109.5^o due to compression by the lone pair electrons.

Oxygen

Bonding in H₂O

The oxygen in H₂O has six valence electrons. After hybridization these six electrons are placed in the four equivalent sp³ hybrid orbitals. The electron configuration of oxygen now has two sp³ hybrid orbitals completely filled with two electrons and two sp³ hybrid orbitals with one unpaired electron each. The filled sp³ hybrid orbitals are considered non-bonding because they are already paired. These electrons will be represented as a two sets of lone pair on the structure of H₂O . The two unpaired electrons in the hybrid orbitals are considered bonding and will overlap with the s orbitals in hydrogen to form O-H sigma bonds. Note! This bonding configuration was predicted by the Lewis structure of H₂O.

The four sp³ hybrid orbitals of oxygen orientate themselves to form a tetrahedral geometry. The two O-H sigma bonds of H₂O are formed by sp³(O)-1s(H) orbital overlap. The two remaining sp³ hybrid orbitals each contain two electrons in the form of a lone pair.

Methanol

The oxygen is sp³hybridized which means that it has four sp³ hybrid orbitals. One of the sp³hybridized orbitals overlap with s orbitals from a hydrogen to form the O-H sigma bonds. One of the sp³hybridized orbitals overlap with an sp³ hybridized orbital from carbon to form the C-O sigma bond. Both the sets of lone pair electrons on the oxygen are contained in the remaining sp³ hybridized orbital. Due to the sp³ hybridization the oxygen has a tetrahedral geometry. However, the H-O-C bond angles are less than the typical 109.5^o due to compression by the lone pair electrons.

Phosphorus

Methyl phosphate

The bond pattern of phosphorus is analogous to nitrogen because they are both in period 15. However, phosphorus can have have expanded octets because it is in the n = 3 row. Typically, phosphorus forms five covalent bonds. 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.

Phosphorus Atoms

• 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

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

Objective

After completing this section, you should be able to apply the concept of hybridization of atoms such as N, O, P and S to explain the structures of simple species containing these atoms.

Key Terms

Make certain that you can define, and use in context, the key term below.

• lone pair electrons

Study Notes

Nitrogen is frequently found in organic compounds. As with carbon atoms, nitrogen atoms can be sp³-, sp²- or sp‑hybridized.

Note that, in this course, the term 'lone pair' is used to describe an unshared pair of electrons.

The valence-bond concept of orbital hybridization can be extrapolated to other atoms including nitrogen, oxygen, phosphorus, and sulfur. In other compounds, covalent bonds that are formed can be described using hybrid orbitals.

Nitrogen

Bonding in NH₃

The nitrogen in NH₃ has five valence electrons. After hybridization these five electrons are placed in the four equivalent sp³ hybrid orbitals. The electron configuration of nitrogen now has one sp³ hybrid orbital completely filled with two electrons and three sp³ hybrid orbitals with one unpaired electron each. The two electrons in the filled sp³ hybrid orbital are considered non-bonding because they are already paired. These electrons will be represented as a lone pair on the structure of NH₃. The three unpaired electrons in the hybrid orbitals are considered bonding and will overlap with the s orbitals in hydrogen to form N-H sigma bonds. Note! This bonding configuration was predicted by the Lewis structure of NH₃.

The four sp³ hybrid orbitals of nitrogen orientate themselves to form a tetrahedral geometry. The three N-H sigma bonds of NH₃ are formed by sp³(N)-1s(H) orbital overlap. The fourth sp³ hybrid orbital contains the two electrons of the lone pair and is not directly involved in bonding.

Phosphorus Atomic Number

Methyl amine

The nitrogen is sp³hybridized which means that it has four sp³ hybrid orbitals. Two of the sp³hybridized orbitals overlap with s orbitals from hydrogens to form the two N-H sigma bonds. One of the sp³hybridized orbitals overlap with an sp³ hybridized orbital from carbon to form the C-N sigma bond. The lone pair electrons on the nitrogen are contained in the last sp³ hybridized orbital. Due to the sp³hybridization the nitrogen has a tetrahedral geometry. However, the H-N-H and H-N-C bonds angles are less than the typical 109.5^o due to compression by the lone pair electrons.

Oxygen

Bonding in H₂O

The oxygen in H₂O has six valence electrons. After hybridization these six electrons are placed in the four equivalent sp³ hybrid orbitals. The electron configuration of oxygen now has two sp³ hybrid orbitals completely filled with two electrons and two sp³ hybrid orbitals with one unpaired electron each. The filled sp³ hybrid orbitals are considered non-bonding because they are already paired. These electrons will be represented as a two sets of lone pair on the structure of H₂O . The two unpaired electrons in the hybrid orbitals are considered bonding and will overlap with the s orbitals in hydrogen to form O-H sigma bonds. Note! This bonding configuration was predicted by the Lewis structure of H₂O.

The four sp³ hybrid orbitals of oxygen orientate themselves to form a tetrahedral geometry. The two O-H sigma bonds of H₂O are formed by sp³(O)-1s(H) orbital overlap. The two remaining sp³ hybrid orbitals each contain two electrons in the form of a lone pair.

Methanol

The oxygen is sp³hybridized which means that it has four sp³ hybrid orbitals. One of the sp³hybridized orbitals overlap with s orbitals from a hydrogen to form the O-H sigma bonds. One of the sp³hybridized orbitals overlap with an sp³ hybridized orbital from carbon to form the C-O sigma bond. Both the sets of lone pair electrons on the oxygen are contained in the remaining sp³ hybridized orbital. Due to the sp³ hybridization the oxygen has a tetrahedral geometry. However, the H-O-C bond angles are less than the typical 109.5^o due to compression by the lone pair electrons.

Phosphorus

Methyl phosphate

The bond pattern of phosphorus is analogous to nitrogen because they are both in period 15. However, phosphorus can have have expanded octets because it is in the n = 3 row. Typically, phosphorus forms five covalent bonds. 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.

Phosphorus Atoms

Sulfur

Methanethiol & Dimethyl Sulfide

Sulfur has a bonding pattern similar to oxygen because they are both in period 16 of the periodic table. Because sulfur is positioned in the third row of the periodic table it has the ability to form an expanded octet and the ability to form more than the typical number of covalent bonds. In biological system, sulfur is typically found in molecules called thiols or sulfides. In a thiol, the sulfur atom is bonded to one hydrogen and one carbon and is analogous to an alcohol O-H bond. In a sulfide, the sulfur is bonded to two carbons. The simplest example of a thiol is methane thiol (CH₃SH) and the simplest example of a sulfide is dimethyl sulfide [(CH₃)₃S]. In both cases the sulfur is sp³hybridized, however the sulfur bond angles are much less than the typical tetrahedral 109.5^o being 96.6^o and 99.1^o respectively.

methanethiol

dimethyl sulfide

Exercises

Best luxury crossover suv 2021 7 seater. 1) Insert the missing lone pairs of electrons in the following molecules, and tell what hybridization you expect for each of the indicated atoms.

a) The oxygen is dimethyl ether:

b) The nitrogen in dimethyl amine:

c) The phosphorus in phosphine:

d) The sulfur in hydrogen sulfide:

Solutions

1)

Phosphorus Atom Diagram

a) sp³ hybridization

b) sp³ hybridization

c) sp³ hybridization

d) sp³ hybridization

Questions

Phosphorus Atom Model

Q1.10.1

Identify geometry and lone pairs on each heteroatom of the molecules given.

Solutions

S1.10.1

Diethyl ether would have two lone pairs of electrons and would have a bent geometry around the oxygen.

Phosphorus Atomic Structure

Dimethyl amine would have one lone pair and would show a pyramidal geometry around the nitrogen.

Phosphorus Atomic Size

Contributors and Attributions

• Dr. Used transmissions. Dietmar Kennepohl FCIC (Professor of Chemistry, Athabasca University)

• Prof. Steven Farmer (Sonoma State University)

• Dr. Krista Cunningham