Phosphorus Electrons

1. The number of electrons in an electrically-neutral atom is the same as the number of protons in the nucleus. Therefore, the number of electrons in neutral atom of Phosphorus is 15. Each electron is influenced by the electric fields produced by the positive nuclear charge and the other (Z – 1) negative electrons in the atom.
2. A neutral phosphorus atom in a grounded state will have six electrons occupying orbitals in s sublevels across its electron shells.
3. Bohr Model of Phosphorus Neon Atom Model, 5th Grade Science Projects, 8th copper bohr diagram wedocable - 28 images - copper element protons. As was mentioned before, a neutral Phosphorus Atom contains five valence electrons in the third shell. You can see a Bohr Model of Phosphorus.
4. Electrons are locked up in covalent bond with valence electrons of four neighboring Si atoms. However, the 5th valence electron of phosphorus atom does not find a binding electron and thus remains free to float. When a voltage is applied across the silicon-phosphorus mixture, free electrons migrate toward the positive voltage end.

1. Phosphorus Electrons Configuration
2. Phosphorus Electron Notation
3. Phosphorus Electrons Gained Or Lost
4. Phosphorus Electron Shell
5. Phosphorus Electrons And Neutrons
6. Phosphorus Electrons Lewis Dot Structure

Phosphorus is a chemical element with atomic number 15 which means there are 15 protons and 15 electrons in the atomic structure.

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Phosphorus (P), nonmetallic chemical element of the nitrogen family (Group 15 [Va] of the periodic table) that at room temperature is a colourless, semitransparent, soft, waxy solid that glows in the dark.

History

Arabian alchemists of the 12th century may have isolated elemental phosphorus by accident, but the records are unclear. Phosphorus appears to have been discovered in 1669 by Hennig Brand, a German merchant whose hobby was alchemy. Brand allowed 50 buckets of urine to stand until they putrified and “bred worms.” He then boiled the urine down to a paste and heated it with sand, thereby distilling elemental phosphorus from the mixture. Brand reported his discovery in a letter to Gottfried Wilhelm Leibniz, and, thereafter, demonstrations of this element and its ability to glow in the dark, or “phosphoresce,” excited public interest. Phosphorus, however, remained a chemical curiosity until about a century later when it proved to be a component of bones. Digestion of bones with nitric or sulfuric acid formed phosphoric acid, from which phosphorus could be distilled by heating with charcoal. In the late 1800s James Burgess Readman of Edinburgh developed an electric furnace method for producing the element from phosphate rock, which is essentially the method employed today.

Occurrence and distribution

Phosphorus is a very widely distributed element—12th most abundant in crustEarth’s , to which it contributes about 0.10 weight percent. Its cosmic abundance is about one atom per 100 atoms of silicon, the standard. Its high chemical reactivity assures that it does not occur in the free state (except in a few meteorites). Phosphorus always occurs as the phosphateion. The principal combined forms in nature are the phosphate salts. About 550 different minerals have been found to contain phosphorus, but, of these, the principal source of phosphorus is the apatite series in which calcium ions exist along with phosphate ions and variable amounts of fluoride, chloride, or hydroxide ions, according to the formula [Ca₁₀(PO₄)₆(F, Cl, or OH)₂]. Other important phosphorus-bearing minerals are wavellite and vivianite. Commonly, such metal atoms as magnesium, manganese, strontium, and lead substitute for calcium in the mineral, and silicate, sulfate, vanadate, and similar anions substitute for phosphate ions. Very large sedimentary deposits of fluoroapatite are found in many parts of Earth. The phosphate of bone and tooth enamel is hydroxyapatite. (The principle of lessening tooth decay by fluoridation depends upon the conversion of hydroxyapatite to the harder, more decay-resistant, fluoroapatite.)

The chief commercial source is phosphorite, or phosphate rock, an impure massive form of carbonate-bearing apatite. Estimates of the total phosphate rock in Earth’s crust average about 65,000,000,000 tons, of which Morocco and Western Sahara contain about 80 percent. This estimate includes only ore that is sufficiently rich in phosphate for conversion to useful products by present methods. Vast quantities of material lower in phosphorus content also exist.

The only naturally occurring isotope of phosphorus is that of mass 31. The other isotopes from mass 24 to mass 46 have been synthesized by appropriate nuclear reactions. All of these are radioactive with relatively short half-lives. The isotope of mass 32 has a half-life of 14.268 days and has proven extremely useful in tracer studies involving the absorption and movement of phosphorus in living organisms.

Commercial production and uses

The principal technique for converting phosphate rock to usable materials involves acidulation of the crushed rock—with either sulfuric or phosphoric acids—to form crude calciumhydrogen phosphates that, being water-soluble, are valuable additions to fertilizer. Most of the output is burned to phosphoric anhydride and subsequently treated with water to form phosphoric acid, H₃PO₄. About 95 percent of the phosphate rock mined in the United States is used to make fertilizer or food supplements for animals. Concerns have arisen about phosphorus use, however. Most of the phosphorus is wasted on its journey from mining to being eaten by humans, and the wasted phosphorus ends up in waterways where it can cause algal blooms. Another concern is that increased phosphorus usage will deplete the nonrenewable supply of phosphate rock.

Only about 5 percent of the phosphorus consumed per year in the United States is used in the elemental form. Pyrotechnic applications of the element include tracers, incendiaries, fireworks, and matches. Some is used as an alloying agent, some is used to kill rodents, and the rest is employed in chemical synthesis. A large amount is converted to sulfides used in matches and in the manufacture of insecticides and oil additives. Most of the remainder is converted to halides or oxides for subsequent use in synthesizing organic phosphorus compounds.

The Chemistry of Nitrogen and Phosphorous

The chemistry of nitrogen is dominated by the ease with which nitrogen atoms form double and triple bonds. A neutral nitrogen atom contains five valence electrons: 2s² 2p³. A nitrogen atom can therefore achieve an octet of valence electrons by sharing three pairs of electrons with another nitrogen atom.

Because the covalent radius of a nitrogen atom is relatively small (only 0.070 nm), nitrogen atoms come close enough together to form very strong bonds. The bond-dissociation enthalpy for the nitrogen-nitrogen triple bond is 946 kJ/mol, almost twice as large as that for an O=O double bond.

The strength of the nitrogen-nitrogen triple bond makes the N₂ molecule very unreactive. N₂ is so inert that lithium is one of the few elements with which it reacts at room temperature.

In spite of the fact that the N₂ molecule is unreactive, compounds containing nitrogen exist for virtually every element in the periodic table except those in Group VIIIA (He, Ne, Ar, and so on). This can be explained in two ways. First, N₂ becomes significantly more reactive as the temperature increases. At high temperatures, nitrogen reacts with hydrogen to form ammonia and with oxygen to form nitrogen oxide.

Second, a number of catalysts found in nature overcome the inertness of N₂ at low temperature.

It is difficult to imagine a living system that does not contain nitrogen, which is an essential component of the proteins, nucleic acids, vitamins, and hormones that make life possible. Animals pick up the nitrogen they need from the plants or other animals in their diet. Plants have to pick up their nitrogen from the soil, or absorb it as N₂ from the atmosphere. The concentration of nitrogen in the soil is fairly small, so the process by which plants reduce N₂ to NH₃ or 'fix' N₂ is extremely important.

Although 200 million tons of NH₃ are produced by nitrogen fixation each year, plants, by themselves, cannot reduce N₂ to NH₃. This reaction is carried out by blue-green algae and bacteria that are associated with certain plants. The best-understood example of nitrogen fixation involves the rhizobium bacteria found in the root nodules of legumes such as clover, peas and beans. These bacteria contain a nitrogenase enzyme, which is capable of the remarkable feat of reducing N₂ from the atmosphere to NH₃ at room temperature.

Ammonia is made on an industrial scale by a process first developed between 1909 and 1913 by Fritz Haber. In the Haber process, a mixture of N₂ and H₂ gas at 200 to 300 atm and 400 to 600^oC is passed over a catalyst of finely divided iron metal.

Almost 20 million tons of NH₃ are produced in the United States each year by this process. About 80% of it, worth more than $2 billion, is used to make fertilizers for plants that can't fix nitrogen from the atmosphere. On the basis of weight, ammonia is the second most important industrial chemical in the United States. (Only sulfuric acid is produced in larger quantities.)

Two-thirds of the ammonia used for fertilizers is converted into solids such as ammonium nitrate, NH₄NO₃; ammonium phosphate, (NH₄)₃PO₄; ammonium sulfate, (NH₄)₂SO₄; and urea, H₂NCONH₂. The other third is applied directly to the soil as anhydrous (literally, 'without water') ammonia. Ammonia is a gas at room temperature. It can be handled as a liquid when dissolved in water to form an aqueous solution. Alternatively, it can be cooled to temperatures below -33^oC, in which case the gas condenses to form the anhydrous liquid, NH₃(l).

The NH₃ produced by the Haber process that is not used as fertilizer is burned in oxygen to generate nitrogen oxide.

Nitrogen oxideor nitric oxide, as it was once known is a colorless gas that reacts rapidly with oxygen to produce nitrogen dioxide, a dark brown gas.

Nitrogen dioxide dissolves in water to give nitric acid and NO, which can be captured and recycled.

Thus, by a three-step process developed by Friedrich Ostwald in 1908, ammonia can be converted into nitric acid.

The Haber process for the synthesis of ammonia combined with the Ostwald process for the conversion of ammonia into nitric acid revolutionized the explosives industry. Nitrates have been important explosives ever since Friar Roger Bacon mixed sulfur, saltpeter, and powdered carbon to make gunpowder in 1245.

Before the Ostwald process was developed the only source of nitrates for use in explosives was naturally occurring minerals such as saltpeter, which is a mixture of NaNO₃ and KNO₃. Once a dependable supply of nitric acid became available from the Ostwald process, a number of nitrates could be made for use as explosives. Combining NH₃ from the Haber process with HNO₃ from the Ostwald process, for example, gives ammonium nitrate, which is both an excellent fertilizer and a cheap, dependable explosive commonly used in blasting powder.

The destructive power of ammonia nitrate is apparent in photographs of the Alfred P. Murrah Federal Building in Oklahoma City, which was destroyed with a bomb made from ammonium nitrate on April 19, 1995.

Nitric acid (HNO₃) and ammonia (NH₃) represent the maximum (+5) and minimum (-3) oxidation numbers for nitrogen. Nitrogen also forms compounds with every oxidation number between these extremes (see table below).

Common Oxidation Numbers for Nitrogen

At about the time that Haber developed the process for making ammonia and Ostwald worked out the process for converting ammonia into nitric acid, Raschig developed a process that used the hypochlorite (OCl^-) ion to oxidize ammonia to produce hydrazine, N₂H₄.

This reaction can be understood by noting that the OCl^- ion is a two-electron oxidizing agent. The loss of a pair of electrons and a pair of H⁺ ions by neighboring NH₃ molecules would produce a pair of highly reactive NH₂ molecules, which would combine to form a hydrazine molecule as shown in the figure below.

Hydrazine is a colorless liquid with a faint odor of ammonia that can be collected when this solution is heated until N₂H₄ distills out of the reaction flask. Many of the physical properties of hydrazine are similar to those of water.

There is a significant difference between the chemical properties of these compounds, however. Hydrazine burns when ignited in air to give nitrogen gas, water vapor, and large amounts of energy.

The principal use of hydrazine is as a rocket fuel. It is second only to liquid hydrogen in terms of the number of kilograms of thrust produced per kilogram of fuel burned. Hydrazine has several advantages over liquid H₂, however. It can be stored at room temperature, whereas liquid hydrogen must be stored at temperatures below -253^oC. Hydrazine is also more dense than liquid H₂ and therefore requires less storage space.

Pure hydrazine is seldom used as a rocket fuel because it freezes at the temperatures encountered in the upper atmosphere. Hydrazine is mixed with N,N-dimethylhydrazine, (CH₃)₂NNH₂, to form a solution that remains a liquid at low temperatures. Mixtures of hydrazine and N,N-dimethylhydrazine were used to fuel the Titan II rockets that carried the Project Gemini spacecraft, and the reaction between hydrazine derivatives and N₂O₄ is still used to fuel the small rocket engines that enable the space shuttle to maneuver in space.

The product of the combustion of hydrazine is unusual. When carbon compounds burn, the carbon is oxidized to CO or CO₂. When sulfur compounds burn, SO₂ is produced. When hydrazine is burned, the product of the reaction is N₂ because of the unusually strong nitrogen-nitrogen triple bond in the N₂ molecule.

Phosphorus Electrons Configuration

Hydrazine reacts with nitrous acid (HNO₂) to form hydrogen azide, HN₃, in which the nitrogen atom formally has an oxidation state of -¹/₃.

Pure hydrogen azide is an extremely dangerous substance. Even dilute solutions should be handled with care because of the risk of explosions. Hydrogen azide is best described as a resonance hybrid of the Lewis structures shown in the figure below. The corresponding azide ion, N₃^-, is a linear molecule, which is a resonance hybrid of three Lewis structures.

Fluorine, oxygen, and chlorine are the only elements more electronegative than nitrogen. As a result, positive oxidation numbers of nitrogen are found in compounds that contain one or more of these elements.

In theory, N₂ could react with F₂ to form a compound with the formula NF₃. In practice, N₂ is too inert to undergo this reaction at room temperature. NF₃ is made by reacting ammonia with F₂ in the presence of a copper metal catalyst.

The HF produced in this reaction combines with ammonia to form ammonium fluoride. The overall stoichiometry for the reaction is therefore written as follows.

The Lewis structure of NF₃ is analogous to the Lewis structure of NH₃, and the two molecules have similar shapes.

Ammonia reacts with chlorine to form NCl₃, which seems at first glance to be closely related to NF₃. But there is a significant difference between these compounds. NF₃ is essentially inert at room temperature, whereas NCl₃ is a shock-sensitive, highly explosive liquid that decomposes to form N₂ and Cl₂.

Ammonia reacts with iodine to form a solid that is a complex between NI₃ and NH₃. This material is the subject of a popular, but dangerous, demonstration in which freshly prepared samples of NI₃ in ammonia are poured onto filter paper, which is allowed to dry on a ring stand. After the ammonia evaporates, the NH₃/NI₃ crystals are touched with a feather attached to a meter stick, resulting in detonation of this shock-sensitive solid, which decomposes to form a mixture of N₂ and I₂.

Lewis structures for seven oxides of nitrogen with oxidation numbers ranging from +1 to +5 are given in the table below.

These compounds all have two things in common: they contain N=O double bonds and they are less stable than their elements in the gas phase, as shown by the enthalpy of formation data in the table below.

Enthalpy of Formation Data for the Oxides of Nitrogen

Dinitrogen oxide, N₂O, which is also known as nitrous oxide, can be prepared by carefully decomposing ammonium nitrate.

Nitrous oxide is a sweet-smelling, colorless gas best known to nonchemists as 'laughing gas.' As early as 1800, Humphry Davy noted that N₂O, inhaled in relatively small amounts, produced a state of apparent intoxication often accompanied by either convulsive laughter or crying. When taken in larger doses, nitrous oxide provides fast and efficient relief from pain. N₂O was therefore used as the first anesthetic. Because large doses are needed to produce anesthesia, and continued exposure to the gas can be fatal, N₂O is used today only for relatively short operations.

Nitrous oxide has several other interesting properties. First, it is highly soluble in cream; for that reason, it is used as the propellant in whipped cream dispensers. Second, although it does not burn by itself, it is better than air at supporting the combustion of other objects. This can be explained by noting that N₂O can decompose to form an atmosphere that is one-third O₂ by volume, whereas normal air is only 21% oxygen by volume.

For many years, the endings -ous and -ic were used to distinguish between the lowest and highest of a pair of oxidation numbers. N₂O is nitrous oxide because the oxidation number of the nitrogen is +1. NO is nitric oxide because the oxidation number of the nitrogen is +2.

Enormous quantities of nitrogen oxide, or nitric oxide, are generated each year by the reaction between the N₂ and O₂ in the atmosphere, catalyzed by a stroke of lightning passing through the atmosphere or by the hot walls of an internal combustion engine.

One of the reasons for lowering the compression ratio of automobile engines in recent years is to decrease the temperature of the combustion reaction, thereby decreasing the amount of NO emitted into the atmosphere.

NO can be prepared in the laboratory by reacting copper metal with dilute nitric acid.

The NO molecule contains an odd number of valence electrons. As a result, it is impossible to write a Lewis structure for this molecule in which all of the electrons are paired (see table of oxides of nitrogen). When NO gas is cooled, pairs of NO molecules combine in a reversible reaction to form a dimer (from the Greek, 'two parts'), with the formula N₂O₂, in which all of the valence electrons are paired, as shown in the table of oxides of nitrogen.

NO reacts rapidly with O₂ to form nitrogen dioxide (once known as nitrogen peroxide), which is a dark brown gas at room temperature.

NO₂ can be prepared in the laboratory by heating certain metal nitrates until they decompose.

It can also be made by reacting copper metal with concentrated nitric acid,

NO₂ also has an odd number of electrons and therefore contains at least one unpaired electron in its Lewis structures. NO₂ dimerizes at low temperatures to form N₂O₄ molecules, in which all the electrons are paired, as shown in the table of oxides of nitrogen.

Mixtures of NO and NO₂ combine when cooled to form dinitrogen trioxide, N₂O₃, which is a blue liquid. The formation of a blue liquid when either NO or NO₂ is cooled therefore implies the presence of at least a small portion of the other oxide because N₂O₂ and N₂O₄ are both colorless.

By carefully removing water from concentrated nitric acid at low temperatures with a dehydrating agent we can form dinitrogen pentoxide.

N₂O₅ is a colorless solid that decomposes in light or on warming to room temperature. As might be expected, N₂O₅ dissolves in water to form nitric acid.

Phosphorus is the first element whose discovery can be traced to a single individual. In 1669, while searching for a way to convert silver into gold, Hennig Brand obtained a white, waxy solid that glowed in the dark and burst spontaneously into flame when exposed to air. Brand made this substance by evaporating the water from urine and allowing the black residue to putrefy for several months. He then mixed this residue with sand, heated this mixture in the presence of a minimum of air, and collected under water the volatile products that distilled out of the reaction flask.

Phosphorus forms a number of compounds that are direct analogs of nitrogen-containing compounds. However, the fact that elemental nitrogen is virtually inert at room temperature, whereas elemental phosphorus can burst spontaneously into flame when exposed to air, shows that there are differences between these elements as well. Phosphorus often forms compounds with the same oxidation numbers as the analogous nitrogen compounds, but with different formulas, as shown in the table below.

Nitrogen and Phosphorus Compounds with the Same Oxidation Numbers but Different Formulas

The same factors that explain the differences between sulfur and oxygen can be used to explain the differences between phosphorus and nitrogen.

1. Nitrogen-nitrogen triple bonds are much stronger than phosphorus-phosphorus triple bonds.

2. P-P single bonds are stronger than N-N single bonds.

3. Phosphorus (EN = 2.19) is much less electronegative than nitrogen (EN = 3.04).

4. Phosphorus can expand its valence shell to hold more than eight electrons, but nitrogen cannot.

Phosphorus Electron Notation

The ratio of the radii of phosphorus and nitrogen atoms is the same as the ratio of the radii of sulfur and oxygen atoms, within experimental error.

As a result, phosphorus-phosphorus triple bonds are much weaker than Nitrogen-nitrogen triple bonds, for the same reason that S=S double bonds are weaker than O=O double bonds phosphorus atoms are too big to come close enough together to form strong bonds.

Each atom in an N₂ molecule completes its octet of valence electrons by sharing three pairs of electrons with a single neighboring atom. Because phosphorus does not form strong multiple bonds with itself, elemental phosphorus consists of tetrahedral P₄ molecules in which each atom forms single bonds with three neighboring atoms, as shown in the figure below.

Phosphorus is a white solid with a waxy appearance, which melts at 44.1^oC and boils at 287^oC. It is made by reducing calcium phosphate with carbon in the presence of silica (sand) at very high temperatures.

White phosphorus is stored under water because the element spontaneously bursts into flame in the presence of oxygen at temperatures only slightly above room temperature. Although phosphorus is insoluble in water, it is very soluble in carbon disulfide. Solutions of P₄ in CS₂ are reasonably stable. As soon as the CS₂ evaporates, however, the phosphorus bursts into flame.

The P-P-P bond angle in a tetrahedral P₄ molecule is only 60^o. This very small angle produces a considerable amount of strain in the P₄ molecule, which can be relieved by breaking one of the P-P bonds. Phosphorus therefore forms other allotropes by opening up the P₄ tetrahedron. When white phosphorus is heated to 300^oC, one bond inside each P₄ tetrahedron is broken, and the P₄ molecules link together to form a polymer (from the Greek pol, 'many,' and meros, 'parts') with the structure shown in the figure below. This allotrope of phosphorus is dark red, and its presence in small traces often gives white phosphorus a light yellow color. Red phosphorus is more dense (2.16 g/cm³) than white phosphorus (1.82 g/cm³) and is much less reactive at normal temperatures.

The size of a phosphorus atom also interferes with its ability to form double bonds to other elements, such as oxygen, nitrogen, and sulfur. As a result, phosphorus tends to form compounds that contain two P-O single bonds where nitrogen would form an N=O double bond. Nitrogen forms the nitrate, NO₃^-, ion, for example, in which it has an oxidation number of +5. When phosphorus forms an ion with the same oxidation number, it is the phosphate, PO₄^3-, ion, as shown in the figure below.

Similarly, nitrogen forms nitric acid, HNO₃, which contains an N=O double bond, whereas phosphorus forms phosphoric acid, H₃PO₄, which contains P-O single bonds, as shown in the figure below.

The difference between the electronegativities of phosphorus and nitrogen (EN = 0.85) is the same as the difference between the electronegativities of sulfur and oxygen (EN = 0.86), within experimental error. Because it is less electronegative, phosphorus is more likely than nitrogen to exhibit positive oxidation numbers. The most important oxidation numbers for phosphorus are -3, +3, and +5 (see table below).

Common Oxidation Numbers of Phosphorus

Because it is more electronegative than most metals, phosphorus reacts with metals at elevated temperatures to form phosphides, in which it has an oxidation number of -3.

Phosphorus Electrons Gained Or Lost

These metal phosphides react with water to produce a poisonous, highly reactive, colorless gas known as phosphine (PH₃), which has the foulest odor the authors have encountered.

Samples of PH₃, the phosphorus analog of ammonia, are often contaminated by traces of P₂H₄, the phosphorus analog of hydrazine. As if the toxicity and odor of PH₃ were not enough, mixtures of PH₃ and P₂H₄ burst spontaneously into flame in the presence of oxygen.

Compounds (such as Ca₃P₂ and PH₃) in which phosphorus has a negative oxidation number are far outnumbered by compounds in which the oxidation number of phosphorus is positive. Phosphorus burns in O₂ to produce P₄O₁₀ in a reaction that gives off extraordinary amounts of energy in the form of heat and light.

When phosphorus burns in the presence of a limited amount of O₂, P₄O₆ is produced.

P₄O₆ consists of a tetrahedron in which an oxygen atom has been inserted into each P-P bond in the P₄ molecule (see figure below). P₄O₁₀ has an analogous structure, with an additional oxygen atom bound to each of the four phosphorus atoms.

P₄O₆ and P₄O₁₀ react with water to form phosphorous acid, H₃PO₃, and phosphoric acid, H₃PO₄, respectively.

Phosphorus Electron Shell

Phosphorus Electrons And Neutrons

P₄O₁₀ has such a high affinity for water that it is commonly used as a dehydrating agent. Phosphorous acid, H₃PO₃, and phosphoric acid, H₃PO₄, are examples of a large class of oxyacids of phosphorus. Lewis structures for some of these oxyacids and their related oxyanions are given in the table below.

The reaction between ammonia and fluorine stops at NF₃ because nitrogen uses the 2s, 2p_x, 2p_y and 2p_z orbitals to hold valence electrons. Nitrogen atoms can therefore hold a maximum of eight valence electrons. Phosphorus, however, has empty 3d atomic orbitals that can be used to expand the valence shell to hold 10 or more electrons. Thus, phosphorus can react with fluorine to form both PF₃ and PF₅. Phosphorus can even form the PF₆^- ion, in which there are 12 valence electrons on the central atom, as shown in the figure below.

Phosphorus Electrons Lewis Dot Structure