NITROGEN

NITROGEN

CAS Number: 7727-37-9
EC Number: 231-783-9

Nitrogen is the chemical element with the symbol N and atomic number 7. 
Nitrogen is a nonmetal and the lightest member of group 15 of the periodic table, often called the pnictogens. 
Nitrogen is a common element in the universe, estimated at seventh in total abundance in the Milky Way and the Solar System. 

At standard temperature and pressure, two atoms of the element bind to form N2, a colorless and odorless diatomic gas. 
N2 forms about 78% of Earth's atmosphere, making Nitrogen the most abundant uncombined element. 

Nitrogen occurs in all organisms, primarily in amino acids (and thus proteins), in the nucleic acids (DNA and RNA) and in the energy transfer molecule adenosine triphosphate. 

The human body contains about 3% nitrogen by mass, the fourth most abundant element in the body after oxygen, carbon, and hydrogen. 
The Nitrogen cycle describes the movement of the element from the air, into the biosphere and organic compounds, then back into the atmosphere.

Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. 
The extremely strong triple bond in elemental nitrogen (N≡N), the second strongest bond in any diatomic molecule after carbon monoxide (CO), dominates nitrogen chemistry. 

This causes difficulty for both organisms and industry in converting N2 into useful compounds, but at the same time it means that burning, exploding, or decomposing nitrogen compounds to form nitrogen gas releases large amounts of often useful energy. 

Nitrogen was first discovered and isolated by Scottish physician Daniel Rutherford in 1772. 
Although Carl Wilhelm Scheele and Henry Cavendish had independently done so at about the same time, Rutherford is generally accorded the credit because his work was published first. 

The name nitrogène was suggested by French chemist Jean-Antoine-Claude Chaptal in 1790 when it was found that nitrogen was present in nitric acid and nitrates. 

Antoine Lavoisier suggested instead the name azote, from the Ancient Greek: ἀζωτικός "no life", as Nitrogen is an asphyxiant gas; this name is used in several languages, including French, Italian, Russian, Romanian, Portuguese and Turkish, and appears in the English names of some nitrogen compounds such as hydrazine, azides and azo compounds.

The wheat sheaf symbol and lightning reflect the importance of nitrogen to living things. 
Nitrogen is important for plant growth and can be ‘fixed’ by lightning or added to soils in fertilisers.
A colourless, odourless gas.

Nitrogen makes up 78% of the air, by volume. 
Nitrogen is obtained by the distillation of liquid air. 

Around 45 million tonnes are extracted each year. 
Nitrogen is found, as compounds, in all living things and hence also in coal and other fossil fuels.

Nitrogen in the form of ammonium chloride, NH4Cl, was known to the alchemists as sal ammonia. 
Nitrogen was manufactured in Egypt by heating a mixture of dung, salt and urine. 
Nitrogen gas itself was obtained in the 1760s by both Henry Cavendish and Joseph Priestley and they did this by removing the oxygen from air. 

They noted Nitrogen extinguished a lighted candle and that a mouse breathing Nitrogen would soon die. 
Neither man deduced that Nitrogen was an element. 
The first person to suggest this was a young student Daniel Rutherford in his doctorate thesis of September 1772 at Edinburgh, Scotland.

Nitrogen is a non-flammable air gas that forms 78% of the earth's atmosphere.
Nitrogen is one of the primary nutrients critical for the survival of all living organisms. 
Nitrogen is a necessary component of many biomolecules, including proteins, DNA, and chlorophyll. 

Although nitrogen is very abundant in the atmosphere as dinitrogen gas (N2), Nitrogen is largely inaccessible in this form to most organisms, making nitrogen a scarce resource and often limiting primary productivity in many ecosystems. 
Only when nitrogen is converted from dinitrogen gas into ammonia (NH3) does Nitrogen become available to primary producers, such as plants.

In addition to N2 and NH3, nitrogen exists in many different forms, including both inorganic (e.g., ammonia, nitrate) and organic (e.g., amino and nucleic acids) forms. 
Thus, nitrogen undergoes many different transformations in the ecosystem, changing from one form to another as organisms use it for growth and, in some cases, energy. 

The major transformations of nitrogen are nitrogen fixation, nitrification, denitrification, anammox, and ammonification. 
The transformation of nitrogen into Nitrogen's many oxidation states is key to productivity in the biosphere and is highly dependent on the activities of a diverse assemblage of microorganisms, such as bacteria, archaea, and fungi.

Since the mid-1900s, humans have been exerting an ever-increasing impact on the global nitrogen cycle. 
Human activities, such as making fertilizers and burning fossil fuels, have significantly altered the amount of fixed nitrogen in the Earth's ecosystems. 

In fact, some predict that by 2030, the amount of nitrogen fixed by human activities will exceed that fixed by microbial processes. 
Increases in available nitrogen can alter ecosystems by increasing primary productivity and impacting carbon storage. 
Because of the importance of nitrogen in all ecosystems and the significant impact from human activities, nitrogen and its transformations have received a great deal of attention from ecologists.

Nitrogen is an element with atomic symbol N, atomic number 7, and atomic weight 14.01.
Nitrogen, refrigerated liquid (cryogenic liquid) appears as colorless odorless liquid. 
Nontoxic.
Nitrogen is a common normally colourless, odourless, tasteless and mostly diatomic non-metal gas. 

Nitrogen has five electrons in its outer shell, so it is trivalent in most compounds.
Nitrogen constitutes 78 percent of Earth's atmosphere and is a constituent of all living tissues. 
Nitrogen is an essential element for life, because Nitrogen is a constituent of DNA and, as such, is part of the genetic code.

Nitrogen molecules occur mainly in air. 
In water and soils nitrogen can be found in nitrates and nitrites. 
All of these substances are a part of the nitrogen cycle, and there are all interconnected.

Humans have changed natural nitrate and nitrite proportions radically, mainly due to the application of nitrate-containing manures. 
Nitrogen is emitted extensively by industrial companies, increasing the nitrate and nitrite supplies in soil and water as a consequence of reactions that take place in the nitrogen cycle. 

Nitrate concentrations in drinking water will greatly increase due to this.
Nitrogen (N), nonmetallic element of Group 15 [Va] of the periodic table. 
Nitrogen is a colourless, odourless, tasteless gas that is the most plentiful element in Earth’s atmosphere and is a constituent of all living matter.

Nitrogen forms many thousands of organic compounds. 
Most of the known varieties may be regarded as derived from ammonia, hydrogen cyanide, cyanogen, and nitrous or nitric acid. 
The amines, amino acids, and amides, for example, are derived from or closely related to ammonia. 

Nitroglycerin and nitrocellulose are esters of nitric acid. 
Nitro compounds are obtained from the reaction (called nitration) between nitric acid and an organic compound. 
Nitrites are derived from nitrous acid (HNO2). 

Nitroso compounds are obtained by the action of nitrous acid on an organic compound. 
Purines and alkaloids are heterocyclic compounds in which nitrogen replaces one or more carbon atoms.
Nitrogen is a colourless, odourless gas, which condenses at −195.8 °C to a colourless, mobile liquid. 

The element exists as N2 molecules, represented as :N:::N:, for which the bond energy of 226 kilocalories per mole is exceeded only by that of carbon monoxide, 256 kilocalories per mole. 
Because of this high bond energy the activation energy for reaction of molecular nitrogen is usually very high, causing nitrogen to be relatively inert to most reagents under ordinary conditions. 

Furthermore, the high stability of the nitrogen molecule contributes significantly to the thermodynamic instability of many nitrogen compounds, in which the bonds, although reasonably strong, are far less so than those in molecular nitrogen. 
For these reasons, elemental nitrogen appears to conceal quite effectively the truly reactive nature of its individual atoms.

A relatively recent and unexpected discovery is that nitrogen molecules are able to serve as ligands in complex coordination compounds. 
The observation that certain solutions of ruthenium complexes can absorb atmospheric nitrogen has led to hope that one day a simpler and better method of nitrogen fixation may be found.

An active form of nitrogen, presumably containing free nitrogen atoms, can be created by passage of nitrogen gas at low pressure through a high-tension electrical discharge. 
Nitrogen glows with a yellow light and is much more reactive than ordinary molecular nitrogen, combining with atomic hydrogen and with sulfur, phosphorus, and various metals, and capable of decomposing nitric oxide, NO, to N2 and O2.

A nitrogen atom has the electronic structure represented by 1s22s22p3. 
The five outer shell electrons screen the nuclear charge quite poorly, with the result that the effective nuclear charge felt at the covalent radius distance is relatively high. 

Thus nitrogen atoms are relatively small in size and high in electronegativity, being intermediate between carbon and oxygen in both of these properties. 
The electronic configuration includes three half-filled outer orbitals, which give the atom the capacity to form three covalent bonds. 

The nitrogen atom should therefore be a very reactive species, combining with most other elements to form stable binary compounds, especially when the other element is sufficiently different in electronegativity to impart substantial polarity to the bonds. 

When the other element is lower in electronegativity than nitrogen, the polarity gives partial negative charge to the nitrogen atom, making its lone-pair electrons available for coordination. 

When the other element is more electronegative, however, the resulting partial positive charge on nitrogen greatly limits the donor properties of the molecule. 

When the bond polarity is low (owing to the electronegativity of the other element being similar to that of nitrogen), multiple bonding is greatly favoured over single bonding. 

If disparity of atomic size prevents such multiple bonding, then the single bond that forms is likely to be relatively weak, and the compound is likely to be unstable with respect to the free elements. 
All of these bonding characteristics of nitrogen are observable in Nitrogen's general chemistry.

Often the percentage of nitrogen in gas mixtures can be determined by measuring the volume after all other components have been absorbed by chemical reagents. 
Decomposition of nitrates by sulfuric acid in the presence of mercury liberates nitric oxide, which can be measured as a gas. 

Nitrogen is released from organic compounds when they are burned over copper oxide, and the free nitrogen can be measured as a gas after other combustion products have been absorbed. 

The well-known Kjeldahl method for determining the nitrogen content of organic compounds involves digestion of the compound with concentrated sulfuric acid (optionally containing mercury, or its oxide, and various salts, depending on the nature of the nitrogen compound). 

In this way, the nitrogen present is converted to ammonium sulfate. 
Addition of an excess of sodium hydroxide releases free ammonia, which is collected in standard acid; the amount of residual acid, which has not reacted with ammonia, is then determined by titration.

Nitrogen is essential to life on Earth. 
Nitrogen is a component of all proteins, and it can be found in all living systems. 
Nitrogen compounds are present in organic materials, foods, fertilizers, explosives and poisons. 

Named after the Greek word nitron, for "native soda," and genes for "forming," nitrogen is the fifth most abundant element in the universe. 
Nitrogen gas constitutes 78 percent of Earth's air, according to the Los Alamos National Laboratory. 
On the other hand, the atmosphere of Mars is only 2.6 percent nitrogen. 

In Nitrogen's gas form, nitrogen is colorless, odorless and generally considered as inert. 
In Nitrogen's liquid form, nitrogen is also colorless and odorless, and looks similar to water, according to Los Alamos.

Nitrogen is an essential element for all forms of life and is the structural component of amino acids from which animal and human tissues, enzymes, and many hormones are made. 
For plant growth, available (fixed) nitrogen is usually the limiting nutrient in natural systems. 

Nitrogen chemistry and overall cycling in the global environment are quite complex due to the number of oxidation states. 
Nitrogen itself has five valence electrons and can be found at oxidation states between −3 and +5. 
Thus, numerous species can form from chemical, biochemical, geochemical, and biogeochemical processes. 

The seventh element of the periodic table between carbon and oxygen is nitrogen.
Nitrogen’s an important part of amino acids.
Around eighty per cent of the Earth’s atmosphere comprises nitrogen gas.

Nitrogen has no colour, mostly diatomic non-metal gas which is odourless and colourless in nature.
Since Nitrogen has five electrons in its outer shell, most of its compounds are trivalent.
Nitrogen is a constituent of all living tissues. 

Since Nitrogen is a component of DNA and part of a genetic code, Nitrogen is an essential element of life.
Nitrogen is found in nitrates and nitrites in soil and water.
All these substances are part of the nitrogen cycle and interconnected. 

Industrial companies emit nitrogen extensively, increasing nitrite and nitrate content in the ground and water, being the consequence of reactions in the nitrogen cycle.
Nitrogen is an essential nutrient for plant growth, development and reproduction. 

Despite nitrogen being one of the most abundant elements on earth, nitrogen deficiency is probably the most common nutritional problem affecting plants worldwide – nitrogen from the atmosphere and earth's crust is not directly available to plants.

Nitrogen is atomic number 7, which means each nitrogen atom has 7 protons. 
Nitrogen's element symbol is N. 
Nitrogen is odorless, tasteless, and colorless gas at room temperature and pressure. 

Nitrogen's atomic weight is 14.0067.
Nitrogen gas (N2) makes up 78.1% of the volume of the Earth's air. 
Nitrogen's the most common uncombined (pure) element on Earth. 

Nitrogen's estimated to be the 5th or 7th most abundant element in the Solar System and Milky Way.
While the gas is common on Earth, Nitrogen's not so abundant on other planets. 
For example, nitrogen gas is found in the atmosphere of Mars at levels of about 2.6 percent.

Nitrogen is a nonmetal. 
Like other elements in this group, Nitrogen is a poor conductor of heat and electricity and lacks metallic luster in solid form.
Nitrogen gas is relatively inert, but soil bacteria can 'fix' nitrogen into a form that plants and animals can use to make amino acids and proteins.

Ammonia (NH3) is the most important commercial compound of nitrogen. 
Nitrogen is produced by the Haber Process. Natural gas (methane, CH4) is reacted with steam to produce carbon dioxide and hydrogen gas (H2) in a two step process. 

Hydrogen gas and nitrogen gas reacted via the Haber Process to produce ammonia. 
This colorless gas with a pungent odor is easily liquefied (in fact, the liquid is used as a nitrogen fertilizer). 

Ammonia is also used in the production of urea, NH2CONH2, which is used as a fertilizer, used in the plastic industry, and used in the livestock industry as a feed supplement. 
Ammonia is often the starting compound for many other nitrogen compounds.

Nitrogen is a diatomic gas which comprises 78 percent of the earth's atmosphere.
In addition to air, nitrogen is found in the protein matter of all life forms, in some natural gas-hydrocarbon deposits, and in many organic and inorganic compounds.

Nontoxic, only slightly soluble in water and most other liquids, poor conductor of heat and electricity, inert. 
However at high temperatures and pressures, Nitrogen will combine with some reactive metals (such as lithium and magnesium) to form nitrides, as well as with some gaseous elements such as hydrogen and oxygen.

Nitrogen gas (N2) makes up 78.1% of the Earth’s air, by volume. 
The atmosphere of Mars, by comparison, is only 2.6% nitrogen. 
From an exhaustible source in our atmosphere, nitrogen gas can be obtained by liquefaction and fractional distillation. 

Nitrogen is found in all living systems as part of the makeup of biological compounds.
The French chemist Antoine Laurent Lavoisier mistakenly named nitrogen azote, meaning without life. 
However, nitrogen compounds are found in foods, organic materials, fertilizers, poisons, and explosives. 

Nitrogen, as a gas is colorless, odorless, and generally considered an inert element. 
As a liquid (boiling point = minus 195.8°C), it is also colorless and odorless, and is similar in appearance to water. Nitrogen gas can be prepared by heating a water solution of ammonium nitrite (NH4NO3).

Nitrogen (pronounced /ˈnaɪtrədʒɨn/) is a chemical element that has the symbol N and atomic number 7 and atomic mass 14.00674 u. 
Elemental nitrogen is a colorless, odorless, tasteless and mostly inert diatomic gas at standard conditions, constituting 78% by volume of Earth's atmosphere.

Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanides, contain nitrogen. 
The extremely strong bond in elemental nitrogen dominates nitrogen chemistry, causing difficulty for both organisms and industry in converting the N2 into useful compounds, and releasing large amounts of energy when these compounds burn or decay back into nitrogen gas.

The element nitrogen was discovered by Daniel Rutherford, a Scottish physician, in 1772. 
Nitrogen occurs in all living organisms. 
Nitrogen is a constituent element of amino acids and thus of proteins, and of nucleic acids (DNA and RNA). 

Nitrogen resides in the chemical structure of almost all neurotransmitters, and is a defining component of alkaloids, biological molecules produced by many organisms.
Naturally occurring nitrate (NO3) concentrations in lakes and streams are typically less than 4 mg/L, and the concentration of nitrite (NO2) is generally much lower. 

The concentration of ammonia (NH3) in natural waters is generally less than 0.1 mg/L.
Nitrogen, like phosphorus, is an important nutrient for plant growth. 
In rivers and lakes, nitrogen can be dissolved in the water, attached to particles floating in the water and found in the bodies of all living organisms.

Forms of nitrogen in freshwater systems include inorganic nitrogen: 
nitrate (NO3), nitrite (NO2), ammonia (NH3) and ammonium (NH4). 

These inorganic forms of nitrogen are the most bioavailable, meaning they are most easily used and taken up by organisms that live in the water. 
Other forms of nitrogen include organic nitrogen. 
Kjeldahl nitrogen includes ammonia, ammonium and organic nitrogen together.

Nitrogen is a colorless, odorless, tasteless, diatomic and generally inert gas at standard temperature and pressure.
At atmospheric pressure, nitrogen is liquid between 63 K and 77 K.
Liquids colder than this are considerably more expensive to make than liquid nitrogen is.

In the natural world, the nitrogen cycle is of crucial importance to living organisms. 
Nitrogen is taken from the atmosphere and converted to nitrates through lightning storms and nitrogen fixing bacteria. 
The nitrates fertilize plant growth where the nitrogen becomes bound in amino acids, DNA and proteins. 

Nitrogen can then be eaten by animals. 
Eventually the nitrogen from the plants and animals returns to the soil and atmosphere and the cycle repeats.
Nitrogen is a chemical element with an atomic number of 7 (it has seven protons in its nucleus). 

Molecular nitrogen (N2) is a very common chemical compound in which two nitrogen atoms are tightly bound together. 
Molecular nitrogen is a colorless, odorless, tasteless, and inert gas at normal temperatures and pressures.
About 78% of Earth's atmosphere is nitrogen. 

The strong triple-bond between the atoms in molecular nitrogen makes this compound difficult to break apart, and thus nearly inert. 
Nitrogen is one of the most important elements in the chemistry of living creatures. 
For example, nitrogen is part of amino acids, the building blocks of proteins. 

The Nitrogen Cycle traces the path of nitrogen, in many different chemical forms, through the environment and living organisms. 
Certain microbes can take gaseous nitrogen from the air and convert it to ammonia, making it available to plants and other organisms in a process called "nitrogen fixation".

Nitrogen (N) is the 7th element on the periodic table. 
Nitrogen is the fifth most abundant element in the universe, and Nitrogen is also fairly common on Earth. 

Nitrogen is a major component of the Earth's atmosphere - about 78% of the atmosphere is nitrogen. 
Through a process of fractional distillation, nitrogen can be obtained from liquefied air.

PROPERTIES of NITROGEN:
-ATOMIC:
A nitrogen atom has seven electrons. 
In the ground state, they are arranged in the electron configuration 1s2 2s2 2p1x 2p1y 2p1z. 

Nitrogen, therefore, has five valence electrons in the 2s and 2p orbitals, three of which (the p-electrons) are unpaired. 
Nitrogen has one of the highest electronegativities among the elements (3.04 on the Pauling scale), exceeded only by chlorine (3.16), oxygen (3.44), and fluorine (3.98). 
(The light noble gases, helium, neon, and argon, would presumably also be more electronegative, and in fact are on the Allen scale.) 

Following periodic trends, Nitrogen's single-bond covalent radius of 71 pm is smaller than those of boron (84 pm) and carbon (76 pm), while Nitrogen is larger than those of oxygen (66 pm) and fluorine (57 pm). 
The nitride anion, N3−, is much larger at 146 pm, similar to that of the oxide (O2−: 140 pm) and fluoride (F−: 133 pm) anions. 

The first three ionisation energies of nitrogen are 1.402, 2.856, and 4.577 MJ·mol−1, and the sum of the fourth and fifth is 16.920 MJ·mol−1. 
Due to these very high figures, nitrogen has no simple cationic chemistry. 

The lack of radial nodes in the 2p subshell is directly responsible for many of the anomalous properties of the first row of the p-block, especially in nitrogen, oxygen, and fluorine. 
The 2p subshell is very small and has a very similar radius to the 2s shell, facilitating orbital hybridisation. 

Nitrogen also results in very large electrostatic forces of attraction between the nucleus and the valence electrons in the 2s and 2p shells, resulting in very high electronegativities. 

Hypervalency is almost unknown in the 2p elements for the same reason, because the high electronegativity makes it difficult for a small nitrogen atom to be a central atom in an electron-rich three-center four-electron bond since it would tend to attract the electrons strongly to itself. 

Thus, despite nitrogen's position at the head of group 15 in the periodic table, Nitrogen's chemistry shows huge differences from that of its heavier congeners phosphorus, arsenic, antimony, and bismuth.
Nitrogen may be usefully compared to its horizontal neighbours carbon and oxygen as well as its vertical neighbours in the pnictogen column, phosphorus, arsenic, antimony, and bismuth. 

Although each period 2 element from lithium to oxygen shows some similarities to the period 3 element in the next group (from magnesium to chlorine; these are known as diagonal relationships), their degree drops off abruptly past the boron–silicon pair. 

The similarities of nitrogen to sulfur are mostly limited to sulfur nitride ring compounds when both elements are the only ones present.
Nitrogen does not share the proclivity of carbon for catenation. 

Like carbon, nitrogen tends to form ionic or metallic compounds with metals. 
Nitrogen forms an extensive series of nitrides with carbon, including those with chain-, graphitic-, and fullerenic-like structures.

Nitrogen resembles oxygen with Nitrogen's high electronegativity and concomitant capability for hydrogen bonding and the ability to form coordination complexes by donating its lone pairs of electrons. 
There are some parallels between the chemistry of ammonia NH3 and water H2O. 

For example, the capacity of both compounds to be protonated to give NH4+ and H3O+ or deprotonated to give NH2− and OH−, with all of these able to be isolated in solid compounds.
Nitrogen shares with both its horizontal neighbours a preference for forming multiple bonds, typically with carbon, oxygen, or other nitrogen atoms, through pπ–pπ interactions. 

Thus, for example, nitrogen occurs as diatomic molecules and therefore has very much lower melting (−210 °C) and boiling points (−196 °C) than the rest of its group, as the N2 molecules are only held together by weak van der Waals interactions and there are very few electrons available to create significant instantaneous dipoles. 

This is not possible for its vertical neighbours; thus, the nitrogen oxides, nitrites, nitrates, nitro-, nitroso-, azo-, and diazo-compounds, azides, cyanates, thiocyanates, and imino-derivatives find no echo with phosphorus, arsenic, antimony, or bismuth. 

By the same token, however, the complexity of the phosphorus oxoacids finds no echo with nitrogen. 
Setting aside their differences, nitrogen and phosphorus form an extensive series of compounds with one another; these have chain, ring, and cage structures.

ISOTOPES of NITROGEN:
Nitrogen has two stable isotopes: 14N and 15N. 
The first is much more common, making up 99.634% of natural nitrogen, and the second (which is slightly heavier) makes up the remaining 0.366%. 

This leads to an atomic weight of around 14.007 u. 
Both of these stable isotopes are produced in the CNO cycle in stars, but 14N is more common as Nitrogen's neutron capture is the rate-limiting step. 

14N is one of the five stable odd–odd nuclides (a nuclide having an odd number of protons and neutrons); the other four are 2H, 6Li, 10B, and 180mTa.
The relative abundance of 14N and 15N is practically constant in the atmosphere but can vary elsewhere, due to natural isotopic fractionation from biological redox reactions and the evaporation of natural ammonia or nitric acid. 

Biologically mediated reactions (e.g., assimilation, nitrification, and denitrification) strongly control nitrogen dynamics in the soil. 
These reactions typically result in 15N enrichment of the substrate and depletion of the product.

The heavy isotope 15N was first discovered by S. M. Naudé in 1929, and soon after heavy isotopes of the neighbouring elements oxygen and carbon were discovered. 
Nitrogen presents one of the lowest thermal neutron capture cross-sections of all isotopes. 

Nitrogen is frequently used in nuclear magnetic resonance (NMR) spectroscopy to determine the structures of nitrogen-containing molecules, due to Nitrogen's fractional nuclear spin of one-half, which offers advantages for NMR such as narrower line width. 

14N, though also theoretically usable, has an integer nuclear spin of one and thus has a quadrupole moment that leads to wider and less useful spectra. 

15N NMR nevertheless has complications not encountered in the more common 1H and 13C NMR spectroscopy. 
The low natural abundance of 15N (0.36%) significantly reduces sensitivity, a problem which is only exacerbated by its low gyromagnetic ratio, (only 10.14% that of 1H). 

As a result, the signal-to-noise ratio for 1H is about 300 times as much as that for 15N at the same magnetic field strength. 
This may be somewhat alleviated by isotopic enrichment of 15N by chemical exchange or fractional distillation. 

15N-enriched compounds have the advantage that under standard conditions, they do not undergo chemical exchange of their nitrogen atoms with atmospheric nitrogen, unlike compounds with labelled hydrogen, carbon, and oxygen isotopes that must be kept away from the atmosphere. 

The 15N:14N ratio is commonly used in stable isotope analysis in the fields of geochemistry, hydrology, paleoclimatology and paleoceanography, where it is called δ15N.

Of the ten other isotopes produced synthetically, ranging from 12N to 23N, 13N has a half-life of ten minutes and the remaining isotopes have half-lives on the order of seconds (16N and 17N) or milliseconds. 
No other nitrogen isotopes are possible as they would fall outside the nuclear drip lines, leaking out a proton or neutron. 

Given the half-life difference, 13N is the most important nitrogen radioisotope, being relatively long-lived enough to use in positron emission tomography (PET), although its half-life is still short and thus Nitrogen must be produced at the venue of the PET, for example in a cyclotron via proton bombardment of 16O producing 13N and an alpha particle.

The radioisotope 16N is the dominant radionuclide in the coolant of pressurised water reactors or boiling water reactors during normal operation. 
Nitrogen is produced from 16O (in water) via an (n,p) reaction, in which the 16O atom captures a neutron and expels a proton. 
Nitrogen has a short half-life of about 7.1 s, but during Nitrogen's decay back to 16O produces high-energy gamma radiation (5 to 7 MeV). 

Because of this, access to the primary coolant piping in a pressurised water reactor must be restricted during reactor power operation. 
Nitrogen is a sensitive and immediate indicator of leaks from the primary coolant system to the secondary steam cycle, and is the primary means of detection for such leaks.

USES and APPLICATIONS of NITROGEN:
-Apart from Nitrogen's use in fertilisers and energy stores, nitrogen is a constituent of organic compounds as diverse as Kevlar used in high-strength fabric and cyanoacrylate used in superglue. 
-Nitrogen is a constituent of every major pharmacological drug class, including antibiotics. 

-Many drugs are mimics or prodrugs of natural nitrogen-containing signal molecules: for example, the organic nitrates nitroglycerin and nitroprusside control blood pressure by metabolizing into nitric oxide. 
-Many notable nitrogen-containing drugs, such as the natural caffeine and morphine or the synthetic amphetamines, act on receptors of animal neurotransmitters.

-Gas:
The applications of nitrogen compounds are naturally extremely widely varied due to the huge size of this class: hence, only applications of pure nitrogen itself will be considered here. 
Two-thirds (2/3) of nitrogen produced by industry is sold as the gas and the remaining one-third (1/3) as the liquid.

-As a modified atmosphere, pure or mixed with carbon dioxide, to nitrogenate and preserve the freshness of packaged or bulk foods (by delaying rancidity and other forms of oxidative damage). 
-Pure nitrogen as food additive is labeled in the European Union with the E number E941.

-In incandescent light bulbs as an inexpensive alternative to argon.
-In fire suppression systems for Information technology (IT) equipment.
-In the manufacture of stainless steel.

-In the case-hardening of steel by nitriding.
-In some aircraft fuel systems to reduce fire hazard (see inerting system).
-To inflate race car and aircraft tires, reducing the problems of inconsistent expansion and contraction caused by moisture and oxygen in natural air.

-Large amounts of ammonia are then used to create fertilizers, explosives and, through a process known as the Ostwald process, nitric acid (HNO3).
-Nitrogen gas is largely inert and is used as a protective shield in the semiconductor industry and during certain types of welding and soldering operations. 

-However, the liquid nitrogen is used to produce fog in cocktails.
-Oil companies use high pressure nitrogen to help force crude oil to the surface.
-Liquid nitrogen is an inexpensive cryogenic liquid used for refrigeration, preservation of biological samples and for low temperature scientific experimentation. 

-Nitrogen gas is used as a nonflammable protective atmosphere. 
-The liquid form of the element is used to remove warts, as a computer coolant, and for cryogenics. 
-Nitrogen is commonly used during sample preparation in chemical analysis. 

-Nitrogen is used to concentrate and reduce the volume of liquid samples. 
-Directing a pressurised stream of nitrogen gas perpendicular to the surface of the liquid causes the solvent to evaporate while leaving the solute(s) and un-evaporated solvent behind.

-Nitrogen can be used as a replacement, or in combination with, carbon dioxide to pressurise kegs of some beers, particularly stouts and British ales, due to the smaller bubbles it produces, which makes the dispensed beer smoother and headier. 
-A pressure-sensitive nitrogen capsule known commonly as a "widget" allows nitrogen-charged beers to be packaged in cans and bottles.

-Nitrogen is part of many important compounds, such as nitrous oxide, nitroglycerin, nitric acid, and ammonia. 
The triple bond nitrogen forms with other nitrogen atoms is extremely strong and releases considerable energy when broken, which is why is so valuable in explosives and also "strong" materials such as Kevlar and cyanoacrylate glue ("super glue").

-Nitrogen tanks are also replacing carbon dioxide as the main power source for paintball guns. 
-Nitrogen must be kept at higher pressure than CO2, making N2 tanks heavier and more expensive.

-Equipment:
Some construction equipment uses pressurized nitrogen gas to help hydraulic system to provide extra power to devices such as hydraulic hammer. 
Nitrogen gas, formed from the decomposition of sodium azide, is used for the inflation of airbags.

-Liquid:
Liquid nitrogen is a cryogenic liquid which looks like water. 
When insulated in proper containers such as dewar flasks, Nitrogen can be transported and stored with a low rate of evaporative loss.

-Like dry ice, the main use of liquid nitrogen is for cooling to low temperatures. 
Nitrogen is used in the cryopreservation of biological materials such as blood and reproductive cells (sperm and eggs). 
-Nitrogen is used in cryotherapy to remove cysts and warts on the skin by freezing them.

-Nitrogen is used in laboratory cold traps, and in cryopumps to obtain lower pressures in vacuum pumped systems. 
-Nitrogen is used to cool heat-sensitive electronics such as infrared detectors and X-ray detectors. 

-Other uses include freeze-grinding and machining materials that are soft or rubbery at room temperature, shrink-fitting and assembling engineering components, and more generally to attain very low temperatures where necessary. 

-Because of Nitrogen's low cost, liquid nitrogen is often used for cooling even when such low temperatures are not strictly necessary, such as refrigeration of food, freeze-branding livestock, freezing pipes to halt flow when valves are not present, and consolidating unstable soil by freezing whenever excavation is going on underneath.

-Nitrogen is important to the chemical industry. 
Nitrogen is used to make fertilisers, nitric acid, nylon, dyes and explosives. 
To make these products, nitrogen must first be reacted with hydrogen to produce ammonia. 

This is done by the Haber process. 
150 million tonnes of ammonia are produced in this way every year.

-Nitrogen gas is also used to provide an unreactive atmosphere. 
Nitrogen is used in this way to preserve foods, and in the electronics industry during the production of transistors and diodes. 

-Large quantities of nitrogen are used in annealing stainless steel and other steel mill products. 
Annealing is a heat treatment that makes steel easier to work.
-Liquid nitrogen is often used as a refrigerant. 

-Nitrogen is used for storing sperm, eggs and other cells for medical research and reproductive technology. 
-Nitrogen is also used to rapidly freeze foods, helping them to maintain moisture, colour, flavour and texture.
-Nitrogen is not reactive and Nitrogen is excellent for blanketing and is often used as purging gas. 

-Nitrogen can be used to remove contaminants from process streams through methods such as stripping and sparging. 
-Due to Nitrogen's properties it can be used for protection of valuable products against harmful contaminants. 
-Nitrogen also enables safe storage, usage of flammable compounds and can help prevent combustible dust explosions.

-Two-thirds of nitrogen produced by industry is sold as the gas and the remaining one-third as the liquid. 
-Nitrogen, The gas is mostly used as an inert atmosphere whenever the oxygen in the air would pose a fire, explosion, or oxidising hazard. 

-Food industry:
Nitrogen gas is also used to provide an unreactive atmosphere. 
Nitrogen is used in this way to preserve foods. 

As a modified atmosphere, pure or mixed with carbon dioxide, to nitrogenate and preserve the freshness of packaged or bulk foods (by delaying rancidity and other forms of oxidative damage like changing colours). 
Pure nitrogen as food additive is labelled in the European Union with the E number E941.

-Light bulbs industry:
Bulbs should not be filled with air since hot tungsten wire will combust in presence of oxygen. 
You can’t maintain vacuum either or external atmospheric pressure will break the glass. 

So, they must be filled with non-reactive gas like nitrogen. 
We can use inert gases like argon or helium instead of Nitrogen, but they are more expensive & rarer than nitrogen. 

-Fire suppression systems:
Fire suppression is achieved by reducing the oxygen concentration where the fire will extinguish, while remaining at a level acceptable for human exposure for a short period of time.

-Stainless steel manufacturing:
There are various instances when nitrogen can be added to steel during steelmaking such as melting, the ladle processing and the casting operations. 
Nitrogen effect on hardness, formability, strain ageing and impact properties.

-Chemical analysis and chemical industry:
Nitrogen is commonly used during sample preparation in chemical analysis. 
Nitrogen is used to concentrate and reduce the volume of liquid samples.

Nitrogen is also important to the chemical industry. 
Nitrogenis used in production of fertilisers, nitric acid, nylon, dyes and explosives.

-Pressurised beer kegs:
Nitrogen can be used as a replacement, or in combination with, carbon dioxide to pressurise kegs of some beers, particularly stouts and British ales, due to the smaller bubbles it produces, which makes the dispensed beer smoother and headier. Nitrogen-charged beers can be packaged in cans and bottles.

-Tire filling systems
Nitrogen is used to inflate race car and aircraft tires, reducing the problems caused by moisture and oxygen in natural air. Nitrogen is less likely to migrate through tire rubber than oxygen, which means that your tire pressures will remain more stable over the long term. 
That means more consistent inflation pressures during a use as the tires heat up.

-Aircraft fuel systems
In some aircraft fuel systems nitrogen is used to reduce fire hazard. 

-Biological role:
Nitrogen is cycled naturally by living organisms through the ‘nitrogen cycle’. 
Nitrogen is taken up by green plants and algae as nitrates, and used to build up the bases needed to construct DNA, RNA and all amino acids. 
Amino acids are the building blocks of proteins.

-Used to freeze foods, to preserve whole blood and other biologicals, and as a coolant.
-Animals obtain their nitrogen by consuming other living things. 
They digest the proteins and DNA into their constituent bases and amino acids, reforming them for their own use.

-Microbes in the soil convert the nitrogen compounds back to nitrates for the plants to re-use. 
The nitrate supply is also replenished by nitrogen-fixing bacteria that ‘fix’ nitrogen directly from the atmosphere.
-Nitrogen is used to produce ammonia (Haber process) and fertilizers, vital for current food production methods. 

-Nitrogen is also used to manufacture nitric acid (Ostwald process).
-In enhanced oil recovery, high pressure nitrogen is used to force crude oil that would otherwise not be recovered out of oil wells. 

-Nitrogen’s inert qualities find use in the chemical and petroleum industries to blanket storage tanks with an inert layer of gas.
-Liquid nitrogen is used as a refrigerant. 

-Superconductors for practical technologies should ideally have no electrical resistance at temperatures higher than 63 K because this temperature is achievable relatively cheaply using liquid nitrogen. 
Lower temperatures come with a much higher price tag.

-While elemental nitrogen is not very reactive, many of nitrogen’s compounds are unstable. 
Oxides naturally form in steel during welding and these weaken the weld. 
Nitrogen can be used to exclude oxygen during welding, resulting in better welds.

-The inert gas is used in a variety of applications including cutting, purging, cooling and freezing.
-Valued for Nitrogen's inert properties in Nitrogen's gaseous form, nitrogen displaces air and therefore, reduces or eliminates the oxidation of materials. 

-Nitrogen is also used as an assist gas for laser cutting. 
-Given the extremely low temperatures of its liquid state, nitrogen is an ideal gas for cryogenic cooling and freezing.
-Nitrogen can be used in virtually any industry to improve yields and optimize performance. 

-Nitrogen enables the safe storage and use of flammables and prevents the explosion of combustibles. 
-In addition, nitrogen improves the quality and shelf-life of air-sensitive materials such as food, pharmaceuticals and electronic products.

-The greatest single commercial use of nitrogen is as a component in the manufacture of ammonia, subsequently used as fertilizer and to produce nitric acid.
-Liquid nitrogen (often referred to as LN2) is used as a refrigerant for freezing and transporting food products, for the preservation of bodies and reproductive cells (sperm and eggs), and for stable storage of biological samples.

-Nitric acid salts include some important compounds, for example potassium nitrate, nitric acid, and ammonium nitrate. 
-Elemental nitrogen can be used as an inert atmosphere for reactions requiring the exclusion of oxygen and moisture. 

-In the liquid state, nitrogen has valuable cryogenic applications; except for the gases hydrogen, methane, carbon monoxide, fluorine, and oxygen, practically all chemical substances have negligible vapour pressures at the boiling point of nitrogen and exist, therefore, as crystalline solids at that temperature.

-In the chemical industry, nitrogen is used as a preventive of oxidation or other deterioration of a product, as an inert diluent of a reactive gas, as a carrier to remove heat or chemicals and as an inhibitor of fire or explosions. 

-In the food industry nitrogen gas is employed to prevent spoilage through oxidation, mold, or insects, and liquid nitrogen is used for freeze drying and for refrigeration systems. 

-In the electrical industry nitrogen is used to prevent oxidation and other chemical reactions, to pressurize cable jackets, and to shield motors. 
-Nitrogen finds application in the metals industry in welding, soldering, and brazing, where it helps prevent oxidation, carburization, and decarburization. 

-As a nonreactive gas, nitrogen is employed to make foamed—or expanded—rubber, plastics, and elastomers, to serve as a propellant gas for aerosol cans, and to pressurize liquid propellants for reaction jets. 

-In medicine rapid freezing with liquid nitrogen may be used to preserve blood, bone marrow, tissue, bacteria, and semen.
-Liquid nitrogen has also proven useful in cryogenic research.

-Large quantities of nitrogen are used together with hydrogen to produce ammonia, NH3, a colourless gas with a pungent, irritating odour. 
-The chief commercial method of synthesizing ammonia is the Haber-Bosch process. 

-Ammonia is one of the two principal nitrogen compounds of commerce; it has numerous uses in the manufacture of other important nitrogen compounds. 
A large portion of commercially synthesized ammonia is converted into nitric acid (HNO3) and nitrates, which are the salts and esters of nitric acid. 

Ammonia is used in the ammonia-soda process (Solvay process) to produce soda ash, Na2CO3. 
Ammonia is also used in the preparation of hydrazine, N2H4, a colourless liquid used as a rocket fuel and in many industrial processes.

-Nitric acid is another popular commercial compound of nitrogen. 
A colourless, highly corrosive liquid, it is much used in the production of fertilizers, dyes, drugs, and explosives. 

-Urea (CH4N2O) is the most common source of nitrogen in fertilizers. 
Ammonium nitrate (NH4NO3), a salt of ammonia and nitric acid, is also used as a nitrogenous component of artificial fertilizers and, combined with fuel oil, as an explosive (ANFO).

-Nitrous oxide, also known as laughing gas, is sometimes used as an anesthetic; when inhaled it produces mild hysteria. Nitric oxide reacts rapidly with oxygen to form brown nitrogen dioxide, an intermediate in the manufacture of nitric acid and a powerful oxidizing agent utilized in chemical processes and rocket fuels.

-Also of some importance are certain nitrides, solids formed by direct combination of metals with nitrogen, usually at elevated temperatures. 
They include hardening agents produced when alloy steels are heated in an atmosphere of ammonia, a process called nitriding. 

Those of boron, titanium, zirconium, and tantalum have special applications. 
One crystalline form of boron nitride (BN), for example, is nearly as hard as diamond and less easily oxidized and so is useful as a high-temperature abrasive.

-Azides, which may be either inorganic or organic, are compounds that contain three nitrogen atoms as a group, represented as (―N3). 
Most azides are unstable and highly sensitive to shock. 

Some of them, such as lead azide, Pb(N3)2, are used in detonators and percussion caps. 
The azides, like the halogen compounds, readily react with other substances by displacement of the so-called azide group and yield many kinds of compounds.

-In fact, about 80 percent of ammonia that is produced is used as fertilizer. 
-Nitrogen is also used as a refrigerant gas; in the manufacture of plastics, textiles, pesticides and dyes; and in cleaning solutions, according to the New York Department of State.

-Nitrogen is used in the manufacture of ammonia, to produce nitric acid and subsequently used as a fertilizer.
-Nitric acid salts include important compounds like potassium nitrate, ammonium nitrate, and nitric acid. 
Nitrated organic compounds such as nitro glycerine are often explosives.

-Liquid nitrogen is utilized as a refrigerant for transporting foodstuff and freezing purposes. 
Preservation of bodies and reproductive cells and stable storage of biological samples also makes use of liquid nitrogen.
-Nitrogen is used in almost all pharmacological drugs, and is found in nitrous oxide - an anesthetic. 

-Nitrogen is used in a variety of ways in both its gaseous and liquid form, however nitrogen is also a major component of the group of pollutants known collectively as nitrogen oxides or NOx.
-Nitrogen gas can be used to manufacture ammonia (NH3), which is used extensively to produce chemical fertilizers.

-As well, liquid nitrogen is available as a relatively inexpensive cryogenic liquid used to preserve biological specimen and conduct low temperature scientific experiments.
-As well, another major use of nitrogen is in the production of ammonia (NH3) in a process known as the Haber process. 
This ammonia is then used to create fertilizers, explosives, and nitric acid.

-Finally, nitrogen gas is inert (meaning that it is hard to make nitrogen have a chemical reaction) and thus Nitrogen is used to create an atmosphere that prevents chemical reactions when producing semiconductors, as well as in some welding and soldering operations.

-Chemistry labs will also use nitrogen gas to prevent chemical reactions with the oxygen in the atmosphere.
-Nitrogen makes up 78 per cent of the Earth’s atmosphere and is a part of all living tissue. 
Nitrogen is a crucial ingredient of life since Nitrogen is a constituent of DNA and as such is part of the genetic code.

-Nitrogen molecules often exist in the soil. 
Nitrogen can be present in nitrates and nitrites in water and in soil. 
These compounds are all part of the nitrogen cycle and both are interconnected.

-Nitrogen is used as an effective way to prevent oxidation and provides a safe, inert atmosphere which “sweeps” off furnace-generated gases. 
-This is also used as a laser cutting assist steam, that facilitates plasma cutting. 

-Oil companies also use high pressure nitrogen to force crude oil out from the ground. 
-Nitrogen is used in a broad variety of applications for upstream and midstream electricity.

WHERE IS NITROGEN FOUND?
Nitrogen is the fifth most abundant element in the universe, making up about 78 per cent of the atmosphere on earth, and contains an estimated 4,000 trillion tons of gas. 
Nitrogen is extracted through a process called fractional distillation from liquefied air.

HOW DO YOU FIX NITROGEN?
A wide range of microorganisms called diazotrophs, including bacteria such as azotobacter, and archaea, naturally conduct nitrogen fixation in the soil. 
Some nitrogen-fixing bacteria, particularly legumes, have symbiotic relationships with certain plant groups.

WHAT WOULD HAPPEN IF NITROGEN-FIXING BACTERIA DID NOT EXIST? 
Bacteria transform airborne nitrogen and carbon dioxide into functional components that can be used as basic building blocks by plants and animals. 
To living organisms, a loss of all microbes would be terrible news that they can not produce or receive such essential nutrients on their own.

HOW DO PLANTS TAKE UP NITROGEN?
In the form of nitrate (NO3−) and ammonium (NH4+), plants absorb nitrogen from the soil. 
Nitrate is typically the predominant type of absorbed nitrogen available in aerobic soils where nitrification can occur.

Nitrogen is a key component of the bodies of living organisms. 
Nitrogen atoms are found in all proteins. 
In nitrogen fixation, bacteria convert into ammonia, a form of nitrogen usable by plants. 

When animals eat the plants, they acquire usable nitrogen compounds.
Nitrogen is everywhere! In fact, gas makes up about 78% of Earth's atmosphere by volume, far surpassing the  we often think of as "air".

But having nitrogen around and being able to make use of it are two different things. 
Your body, and the bodies of other plants and animals, have no good way to convert  into a usable form. 
We animals—and our plant compatriots—just don't have the right enzymes to capture, or fix, atmospheric nitrogen.

WHERE DOES THAT NITROGEN COME FROM? 
In the natural world, Nitrogen comes from bacteria!
Bacteria play a key role in the nitrogen cycle.

Nitrogen enters the living world by way of bacteria and other single-celled prokaryotes, which convert atmospheric nitrogen into biologically usable forms in a process called nitrogen fixation. 

Some species of nitrogen-fixing bacteria are free-living in soil or water, while others are beneficial symbionts that live inside of plants.

Nitrogen-fixing microorganisms capture atmospheric nitrogen by converting Nitrogen to ammonia which can be taken up by plants and used to make organic molecules. 
The nitrogen-containing molecules are passed to animals when the plants are eaten. 

They may be incorporated into the animal's body or broken down and excreted as waste, such as the urea found in urine.
Nitrogen doesn't remain forever in the bodies of living organisms. 
Instead, Nitrogen's converted from organic nitrogen back into gas by bacteria. 

This process often involves several steps in terrestrial—land—ecosystems. 
Nitrogenous compounds from dead organisms or wastes are converted into ammonia by bacteria, and the ammonia is converted into nitrites and nitrates. 
In the end, the nitrates are made into gas by denitrifying prokaryotes.

Nitrogen cycling in marine ecosystems:
So far, we’ve focused on the natural nitrogen cycle as it occurs in terrestrial ecosystems. 

However, generally similar steps occur in the marine nitrogen cycle. 
There, the ammonification, nitrification, and denitrification processes are performed by marine bacteria and archaea.

Some nitrogen-containing compounds fall to the ocean floor as sediment. 
Over long periods of time, the sediments get compressed and form sedimentary rock. 
Eventually, geological uplift may move the sedimentary rock to land. 

In the past, scientists did not think that this nitrogen-rich sedimentary rock was an important nitrogen source for terrestrial ecosystems. 
However, a new study suggests that it may actually be quite important—the nitrogen is released gradually to plants as the rock wears away, or weathers.

Nitrogen as a limiting nutrient:
In natural ecosystems, many processes, such as primary production and decomposition, are limited by the available supply of nitrogen. 
In other words, nitrogen is often the limiting nutrient, the nutrient that's in shortest supply and thus limits the growth of organisms or populations.

WHAT IS THE NITROGEN CYCLE AND WHY IS IT KEY TO LIFE?
Nitrogen, the most abundant element in our atmosphere, is crucial to life. 
Nitrogen is found in soils and plants, in the water we drink, and in the air we breathe. 

Nitrogen is also essential to life: a key building block of DNA, which determines our genetics, is essential to plant growth, and therefore necessary for the food we grow. 
But as with everything, balance is key: too little nitrogen and plants cannot thrive, leading to low crop yields. 

Plants that do not have enough nitrogen become yellowish and do not grow well and can have smaller flowers and fruits. 
Understanding the Nitrogen Cycle—how nitrogen moves from the atmosphere to earth, through soils and back to the atmosphere in an endless Cycle—can help us grow healthy crops and protect our environment.

WHY IS NITROGEN IMPORTANT?
The delicate balance of substances that is important for maintaining life is an important area of research, and the balance of nitrogen in the environment is no exception. 
When plants lack nitrogen, they become yellowed, with stunted growth, and produce smaller fruits and flowers. 

Farmers may add fertilizers containing nitrogen to their crops, to increase crop growth. 
Without nitrogen fertilizers, scientists estimate that we would lose up to one third of the crops we rely on for food and other types of agriculture. 

NITROGEN IS KEY TO LIFE!
Nitrogen is a key element in the nucleic acids.

DNA Deoxyribonucleic acid, a self-replicating material which is present in nearly all living organisms as the main component of chromosomes, and carrier of genetic information.

RNA Ribonucleic acid, a nucleic acid present in all living cells, acts as a messenger carrying instructions from DNA which are the most important of all biological molecules and crucial for all living things. 
DNA carries the genetic information, which means the instructions for how to make up a life form. 

When plants do not get enough nitrogen, they are unable to produce amino acids (substances that contain nitrogen and hydrogen and make up many of living cells, muscles and tissue). 

Without amino acids, plants cannot make the special proteins that the plant cells need to grow. 
Without enough nitrogen, plant growth is affected negatively. 

WHAT EXACTLY IS THE NITROGEN CYCLE?
The nitrogen cycle is a repeating cycle of processes during which nitrogen moves through both living and non-living things: the atmosphere, soil, water, plants, animals and bacteria.

Microscopic living organisms that usually contain only one cell and are found everywhere. 
Bacteria can cause decomposition or breaking down, of organic material in soils.
In order to move through the different parts of the cycle, nitrogen must change forms. 

In the atmosphere, nitrogen exists as a gas (N2), but in the soils it exists as nitrogen oxide, NO, and nitrogen dioxide, NO2, and when used as a fertilizer, can be found in other forms, such as ammonia, NH3, which can be processed even further into a different fertilizer, ammonium nitrate, or NH4NO3.

There are five stages in the nitrogen cycle, and we will now discuss each of them in turn: fixation or volatilization, mineralization, nitrification, immobilization, and denitrification. 

In this image, microbes in the soil turn nitrogen gas (N2) into what is called volatile ammonia (NH3), so the fixation process is called volatilization. 

Leaching: 
When a mineral or chemical (such as nitrate, or NO3) drains away from soil or other ground material and leaks into surrounding area is where certain forms of nitrogen (such as nitrate, or NO3) becomes dissolved in water.

Stage 1: NITROGEN FIXATION OR VOLATILIZATION
In this stage, nitrogen moves from the atmosphere into the soil. 
Earth’s atmosphere contains a huge pool of nitrogen gas (N2). 

But this nitrogen is “unavailable” to plants, because the gaseous form cannot be used directly by plants without undergoing a transformation. 
To be used by plants, the N2 must be transformed through a process called nitrogen fixation. 
Fixation converts nitrogen in the atmosphere into forms that plants can absorb through their root systems.

A small amount of nitrogen can be fixed when lightning provides the energy needed for N2 to react with oxygen, producing nitrogen oxide, NO, and nitrogen dioxide, NO2. 
These forms of nitrogen then enter soils through rain or snow. 

Nitrogen can also be fixed through the industrial process that creates fertilizer. 
This form of fixing occurs under high heat and pressure, during which atmospheric nitrogen and hydrogen are combined to form ammonia (NH3), which may then be processed further, to produce ammonium nitrate (NH4NO3), a form of nitrogen that can be added to soils and used by plants.

Most nitrogen fixation occurs naturally, in the soil, by bacteria. 
Some bacteria attach to plant roots and have a symbiotic (beneficial for both the plant and the bacteria) relationship with the plant. 

The bacteria get energy through photosynthesis and, in return, they fix nitrogen into a form the plant needs. 
The fixed nitrogen is then carried to other parts of the plant and is used to form plant tissues, so the plant can grow. 

Other bacteria live freely in soils or water and can fix nitrogen without this symbiotic relationship. 
These bacteria can also create forms of nitrogen that can be used by organisms.

Stage 2: MINERALIZATION
This stage takes place in the soil. 
Nitrogen moves from organic materials, such as manure or plant materials to an inorganic form of nitrogen that plants can use. 
Eventually, the plant’s nutrients are used up and the plant dies and decomposes. 

This becomes important in the second stage of the nitrogen cycle. 
Mineralization happens when microbes act on organic material, such as animal manure or decomposing plant or animal material and begin to convert it to a form of nitrogen that can be used by plants. 

All plants under cultivation, except legumes.
A member of the pea family: beans, lentils, soybeans, peanuts and peas, are plants with seed pods that split in half. 
(plants with seed pods that split in half, such as lentils, beans, peas or peanuts) get the nitrogen they require through the soil. 
Legumes get nitrogen through fixation that occurs in their root nodules, as described above.

The first form of nitrogen produced by the process of mineralization is ammonia, NH3. 
The NH3 in the soil then reacts with water to form ammonium, NH4. 
This ammonium is held in the soils and is available for use by plants that do not get nitrogen through the symbiotic nitrogen fixing relationship described above.

Stage 3: NITRIFICATION
The third stage, nitrification, also occurs in soils. 
During nitrification the ammonia in the soils, produced during mineralization, is converted into compounds called nitrites, NO2−, and nitrates, NO3−. 

Nitrates can be used by plants and animals that consume the plants. 
Some bacteria in the soil can turn ammonia into nitrites. 
Although nitrite is not usable by plants and animals directly, other bacteria can change nitrites into nitrates—a form that is usable by plants and animals. 

This reaction provides energy for the bacteria engaged in this process. 
The bacteria that we are talking about are called nitrosomonas and nitrobacter. 
Nitrobacter turns nitrites into nitrates; nitrosomonas transform ammonia to nitrites. 

Both kinds of bacteria can act only in the presence of oxygen, O2. 
The process of nitrification is important to plants, as it produces an extra stash of available nitrogen that can be absorbed by the plants through their root systems.

Stage 4: IMMOBILIZATION
The fourth stage of the nitrogen cycle is immobilization, sometimes described as the reverse of mineralization. 
These two processes together control the amount of nitrogen in soils. 
Just like plants, microorganisms.

An organism, or living thing, that is too tiny to be seen without a microscope, such as a bacterium. 

Living in the soil require nitrogen as an energy source. 
These soil microorganisms pull nitrogen from the soil when the residues of decomposing plants do not contain enough nitrogen. 

When microorganisms take in ammonium (NH4+) and nitrate (NO3−), these forms of nitrogen are no longer available to the plants and may cause nitrogen deficiency, or a lack of nitrogen. 

Immobilization, therefore, ties up nitrogen in microorganisms. 
However, immobilization is important because it helps control and balance the amount of nitrogen in the soils by tying it up, or immobilizing the nitrogen, in microorganisms.

Stage 5: DENITRIFICATION
In the fifth stage of the nitrogen cycle, nitrogen returns to the air as nitrates are converted to atmospheric nitrogen (N2) by bacteria through the process we call denitrification. 

This results in an overall loss of nitrogen from soils, as the gaseous form of nitrogen moves into the atmosphere, back where we began our story.

NITROGEN IS CRUCIAL FOR LIFE!
The cycling of nitrogen through the ecosystem is crucial for maintaining productive and healthy ecosystems with neither too much nor too little nitrogen. 
Plant production and biomass (living material) are limited by the availability of nitrogen. 

Understanding how the plant-soil nitrogen cycle works can help us make better decisions about what crops to grow and where to grow them, so we have an adequate supply of food. 
Knowledge of the nitrogen cycle can also help us reduce pollution caused by adding too much fertilizer to soils. 

Certain plants can uptake more nitrogen or other nutrients, such as phosphorous, another fertilizer, and can even be used as a “buffer,” or filter, to prevent excessive fertilizer from entering waterways. 

As you have seen, not enough nitrogen in the soils leaves plants hungry, while too much of a good thing can be bad. 
Farmers and communities need to work to improve the uptake of added nutrients by crops and treat animal manure waste properly. 

We also need to protect the natural plant buffer zones that can take up nitrogen runoff before it reaches water bodies. 
But, our current patterns of clearing trees to build roads and other construction worsen this problem, because there are fewer plants left to uptake excess nutrients. 

We need to do further research to determine which plant species are best to grow in coastal areas to take up excess nitrogen. We also need to find other ways to fix or avoid the problem of excess nitrogen spilling over into aquatic ecosystems. 

By working toward a more complete understanding of the nitrogen cycle and other cycles at play in Earth’s interconnected natural systems, we can better understand how to better protect Earth’s precious natural resources.

Some nitrogen-fixing organisms are free-living while others are symbiotic nitrogen-fixers, which require a close association with a host to carry out the process. 
Most of the symbiotic associations are very specific and have complex mechanisms that help to maintain the symbiosis. 

For example, root exudates from legume plants (e.g., peas, clover, soybeans) serve as a signal to certain species of Rhizobium, which are nitrogen-fixing bacteria. 

This signal attracts the bacteria to the roots, and a very complex series of events then occurs to initiate uptake of the bacteria into the root and trigger the process of nitrogen fixation in nodules that form on the roots.

Some of these bacteria are aerobic, others are anaerobic; some are phototrophic, others are chemotrophic (i.e., they use chemicals as their energy source instead of light). 

Although there is great physiological and phylogenetic diversity among the organisms that carry out nitrogen fixation, they all have a similar enzyme complex called nitrogenase that catalyzes the reduction of N2 to NH3 (ammonia), which can be used as a genetic marker to identify the potential for nitrogen fixation. 

One of the characteristics of nitrogenase is that the enzyme complex is very sensitive to oxygen and is deactivated in its presence. 
This presents an interesting dilemma for aerobic nitrogen-fixers and particularly for aerobic nitrogen-fixers that are also photosynthetic since they actually produce oxygen. 
Over time, nitrogen-fixers have evolved different ways to protect their nitrogenase from oxygen. 

For example, some cyanobacteria have structures called heterocysts that provide a low-oxygen environment for the enzyme and serves as the site where all the nitrogen fixation occurs in these organisms. 
Other photosynthetic nitrogen-fixers fix nitrogen only at night when their photosystems are dormant and are not producing oxygen.

Genes for nitrogenase are globally distributed and have been found in many aerobic habitats (e.g., oceans, lakes, soils) and also in habitats that may be anaerobic or microaerophilic (e.g., termite guts, sediments, hypersaline lakes, microbial mats, planktonic crustaceans). 

The broad distribution of nitrogen-fixing genes suggests that nitrogen-fixing organisms display a very broad range of environmental conditions, as might be expected for a process that is critical to the survival of all life on Earth.

NITRIFICATION:
Nitrification is the process that converts ammonia to nitrite and then to nitrate and is another important step in the global nitrogen cycle. 
Most nitrification occurs aerobically and is carried out exclusively by prokaryotes. 

There are two distinct steps of nitrification that are carried out by distinct types of microorganisms. 
The first step is the oxidation of ammonia to nitrite, which is carried out by microbes known as ammonia-oxidizers. 

Aerobic ammonia oxidizers convert ammonia to nitrite via the intermediate hydroxylamine, a process that requires two different enzymes, ammonia monooxygenase and hydroxylamine oxidoreductase. 

The process generates a very small amount of energy relative to many other types of metabolism; as a result, nitrosofiers are notoriously very slow growers. 
Additionally, aerobic ammonia oxidizers are also autotrophs, fixing carbon dioxide to produce organic carbon, much like photosynthetic organisms, but using ammonia as the energy source instead of light.

Unlike nitrogen fixation that is carried out by many different kinds of microbes, ammonia oxidation is less broadly distributed among prokaryotes. 
Until recently, it was thought that all ammonia oxidation was carried out by only a few types of bacteria in the genera Nitrosomonas, Nitrosospira, and Nitrosococcus. 

However, in 2005 an archaeon was discovered that could also oxidize ammonia. 
Since their discovery, ammonia-oxidizing Archaea have often been found to outnumber the ammonia-oxidizing Bacteria in many habitats. 

In the past several years, ammonia-oxidizing Archaea have been found to be abundant in oceans, soils, and salt marshes, suggesting an important role in the nitrogen cycle for these newly-discovered organisms. 

Currently, only one ammonia-oxidizing archaeon has been grown in pure culture, Nitrosopumilus maritimus, so our understanding of their physiological diversity is limited.

The second step in nitrification is the oxidation of nitrite (NO2-) to nitrate (NO3-). 
This step is carried out by a completely separate group of prokaryotes, known as nitrite-oxidizing Bacteria. 
Some of the genera involved in nitrite oxidation include Nitrospira, Nitrobacter, Nitrococcus, and Nitrospina. 

Similar to ammonia oxidizers, the energy generated from the oxidation of nitrite to nitrate is very small, and thus growth yields are very low. 

In fact, ammonia- and nitrite-oxidizers must oxidize many molecules of ammonia or nitrite in order to fix a single molecule of CO2. 
For complete nitrification, both ammonia oxidation and nitrite oxidation must occur.

Ammonia-oxidizers and nitrite-oxidizers are ubiquitous in aerobic environments. 
They have been extensively studied in natural environments such as soils, estuaries, lakes, and open-ocean environments. 

However, ammonia- and nitrite-oxidizers also play a very important role in wastewater treatment facilities by removing potentially harmful levels of ammonium that could lead to the pollution of the receiving waters. 

Much research has focused on how to maintain stable populations of these important microbes in wastewater treatment plants. Additionally, ammonia- and nitrite-oxidizers help to maintain healthy aquaria by facilitating the removal of potentially toxic ammonium excreted in fish urine.

ANAMMOX:
Traditionally, all nitrification was thought to be carried out under aerobic conditions, but recently a new type of ammonia oxidation occurring under anoxic conditions was discovered. 

Anammox (anaerobic ammonia oxidation) is carried out by prokaryotes belonging to the Planctomycetes phylum of Bacteria. 
The first described anammox bacterium was Brocadia anammoxidans. 
Anammox bacteria oxidize ammonia by using nitrite as the electron acceptor to produce gaseous nitrogen. 

Anammox bacteria were first discovered in anoxic bioreactors of wasterwater treatment plants but have since been found in a variety of aquatic systems, including low-oxygen zones of the ocean, coastal and estuarine sediments, mangroves, and freshwater lakes. 

In some areas of the ocean, the anammox process is considered to be responsible for a significant loss of nitrogen. 
Whether anammox or denitrification is responsible for most nitrogen loss in the ocean, it is clear that anammox represents an important process in the global nitrogen cycle.

DENITRIFICATION:
Denitrification is the process that converts nitrate to nitrogen gas, thus removing bioavailable nitrogen and returning it to the atmosphere. 
Dinitrogen gas (N2) is the ultimate end product of denitrification, but other intermediate gaseous forms of nitrogen exist. 

Unlike nitrification, denitrification is an anaerobic process, occurring mostly in soils and sediments and anoxic zones in lakes and oceans. 

Similar to nitrogen fixation, denitrification is carried out by a diverse group of prokaryotes, and there is recent evidence that some eukaryotes are also capable of denitrification. 

Some denitrifying bacteria include species in the genera Bacillus, Paracoccus, and Pseudomonas. 
Denitrifiers are chemoorganotrophs and thus must also be supplied with some form of organic carbon.

Denitrification is important in that it removes fixed nitrogen (i.e., nitrate) from the ecosystem and returns it to the atmosphere in a biologically inert form (N2). 
This is particularly important in agriculture where the loss of nitrates in fertilizer is detrimental and costly. 

However, denitrification in wastewater treatment plays a very beneficial role by removing unwanted nitrates from the wastewater effluent, thereby reducing the chances that the water discharged from the treatment plants will cause undesirable consequences (e.g., algal blooms).

AMMONIFICATION:
When an organism excretes waste or dies, the nitrogen in its tissues is in the form of organic nitrogen (e.g. amino acids, DNA).

Various fungi and prokaryotes then decompose the tissue and release inorganic nitrogen back into the ecosystem as ammonia in the process known as ammonification. 
The ammonia then becomes available for uptake by plants and other microorganisms for growth.

The nitrogen cycle, in which atmospheric nitrogen is converted into different nitrogenous compounds, is one the most crucial natural processes to sustain living organisms. 
During the cycle, certain bacteria ‘fix’ atmospheric nitrogen into ammonia, which plants need in order to grow. 

Other bacteria convert the ammonia into amino acids and proteins. 
Animals eat the plants and consume the protein. 

Nitrogen compounds return to the soil through animal and plant waste. 
Bacteria convert the waste nitrogen back to nitrogen gas, which returns to the atmosphere.

In an effort to make crops grow faster, people use nitrogen in fertilisers. 
The nitrogen cycle refers to the movement of nitrogen within and between the atmosphere, biosphere, hydrosphere and geosphere. 

The nitrogen cycle matters because nitrogen is an essential nutrient for sustaining life on Earth. 
Nitrogen is a core component of amino acids, which are the building blocks of proteins, and of nucleic acids, which are the building blocks of genetic material (RNA and DNA). 

When other resources such as light and water are abundant, ecosystem productivity and biomass is often limited by the amount of available nitrogen. 
This is the primary reason why nitrogen is an essential part of fertilizers used to enhance soil quality for agricultural activities.

Nitrogen cycles through both the abiotic and biotic parts of the Earth system. 
The largest reservoir of nitrogen is found in the atmosphere, mostly as nitrogen gas (N2). 

Nitrogen gas makes up 78% of the air we breathe. 
Most nitrogen enters ecosystems via certain kinds of bacteria in soil and plant roots that convert nitrogen gas into ammonia (NH3). 

This process is called nitrogen fixation. 
A very small amount of nitrogen is fixed via lightning interacting with the air. 
Once nitrogen is fixed, other types of bacteria convert ammonia to nitrate (NO3‑) and nitrite (NO2–), which can then be used by other bacteria and plants. 

Consumers (herbivores and predators) get nitrogen compounds from the plants and animals they eat. 
Nitrogen returns to the soil when organisms release waste or die and are decomposed by bacteria and fungi. 

Nitrogen is released back to the atmosphere by bacteria get their energy by breaking down nitrate and nitrite into nitrogen gas (also called denitrification).

FERTILIZER COMPONENT
Nitrogen was discovered in 1772 by chemist and physician Daniel Rutherford, when he removed oxygen and carbon dioxide from air, demonstrating that the residual gas would not support living organisms or combustion, according to the Los Alamos National Laboratory. 

Other scientists, including Carl Wilhelm Scheele and Joseph Priestly, were working on the same problem, and called nitrogen "burnt" air, or air without oxygen. 
In 1786, Antoine Laurent de Lavoisier, called nitrogen "azote," which means "lifeless." 
This was based on the observation that Nitrogen is the part of air cannot support life on Nitrogen's own. 

One of the most important nitrogen compounds is ammonia (NH3), which can be produced in the so-called Haber-Bosch process, in which nitrogen is reacted with hydrogen. 
The colorless ammonia gas with a pungent smell can be easily liquefied into a nitrogen fertilizer. 

THE NITROGEN CYCLE
The nitrogen cycle, in which atmospheric nitrogen is converted into different organic compounds, is one the most crucial natural processes to sustain living organisms. 
During the cycle, bacteria in the soil process or "fix" atmospheric nitrogen into ammonia, which plants need in order to grow. 

Other bacteria convert the ammonia into amino acids and proteins. 
Then animals eat the plants and consume the protein. 

Nitrogen compounds return to the soil through animal waste. 
Bacteria convert the waste nitrogen back to nitrogen gas, which returns to the atmosphere. 

A cycle is a sequence of events or steps that repeats itself regularly. 
In the nitrogen cycle, nitrogen moves from the soil to plants and then to animals and finally back to the soil. 
When Nitrogen returns to the soil from a decaying plant Nitrogen can be used again by another plant.

Nitrogen Cycle has five general steps:
*Nitrogen fixation
*Nitrification
*Denitrification
*Nitrogen assimilation
*Ammonification.

Over many years the actions of people began changing how nitrogen cycled through nature. 
This changed the amount of nitrogen found in living organisms and in the air, soil, and water. 
The balance of nature was upset. 
By understanding how the nitrogen cycle works people can change their actions and protect the environment.

Nitrogen can go through many transformations in the soil. 
These transformations are often grouped into a system called the nitrogen cycle, which can be presented in varying degrees of complexity. 

The nitrogen cycle is appropriate for understanding nutrient and fertilizer management. 
Because microorganisms are responsible for most of these processes, they occur very slowly, if at all, when soil temperatures are below 50° F, but their rates increase rapidly as soils become warmer.

The heart of the nitrogen cycle is the conversion of inorganic to organic nitrogen, and vice versa. 
As microorganisms grow, they remove H₄⁺ and NO₃⁻ from the soil’s inorganic, available nitrogen pool, converting it to organic nitrogen in a process called immobilization. 

When these organisms die and are decomposed by others, excess NH₄⁺ can be released back to the inorganic pool in a process called mineralization. 
Nitrogen can also be mineralized when microorganisms decompose a material containing more nitrogen than they can use at one time, materials such as legume residues or manures. 

Immobilization and mineralization are conducted by most microorganisms, and are most rapid when soils are warm and moist, but not saturated with water. 
The quantity of inorganic nitrogen available for crop use often depends on the amount of mineralization occurring and the balance between mineralization and immobilization.

Ammonium ions (NH₄⁺) not immobilized or taken up quickly by higher plants are usually converted rapidly to NO₃⁻ ions by a process called nitrification. 

This is a two-step process, during which bacteria called Nitrosomonas convert NH₄⁺ to nitrite (NO₂⁻), and then other bacteria, Nitrobacter, convert the NO₂⁻ to NO₃⁻. 

This process requires a well-aerated soil and occurs rapidly enough that one usually finds mostly NO₃⁻ rather than NH₄⁺ in soils during the growing season.

The nitrogen cycle contains several routes by which plant-available nitrogen can be lost from the soil. 
Nitrate-nitrogen is usually more subject to loss than is ammonium-nitrogen. 
Significant loss mechanisms include leaching, denitrification, volatilization and crop removal.

The nitrate form of nitrogen is so soluble that it leaches easily when excess water percolates through the soil. 
This can be a major loss mechanism in coarse-textured soils where water percolates freely, but is less of a problem in finer-textured, more impermeable soils, where percolation is very slow.

These latter soils tend to become saturated easily, and when microorganisms exhaust the free oxygen supply in the wet soil, some obtain it by decomposing NO₃⁻. 

In this process, called denitrification, NO₃⁻ is converted to gaseous oxides of nitrogen or to N₂ gas, both unavailable to plants. 
Denitrification can cause major losses of nitrogen when soils are warm and remain saturated for more than a few days.

Losses of NH₄⁺ nitrogen are less common and occur mainly by volatilization. 
Ammonium ions are basically anhydrous ammonia (NH₃) molecules with an extra hydrogen ion (H⁺) attached. 

When this extra H⁺ is removed from the NH₄ ion by another ion such as hydroxyl (OH⁻), the resulting NH₃ molecule can evaporate, or volatilize from the soil. 
This mechanism is most important in high-pH soils that contain large quantities of OH⁻ ions.

Crop removal represents a loss because nitrogen in the harvested portions of the crop plant is removed from the field completely. 
The nitrogen in crop residues is recycled back into the system, and is better thought of as immobilized rather than removed. 
Much is eventually mineralized and may be reutilized by a crop.

PLANT NITROGEN NEEDS AND UPTAKE
Plants absorb nitrogen from the soil as both NH₄⁺ and NO₃⁻ ions, but because nitrification is so pervasive in agricultural soils, most of the nitrogen is taken up as nitrate. 

Nitrate moves freely toward plant roots as they absorb water. 
Once inside the plant, NO₃⁻ is reduced to an NH₂ form and is assimilated to produce more complex compounds. 

Because plants require very large quantities of nitrogen, an extensive root system is essential to allowing unrestricted uptake. 

Plants with roots restricted by compaction may show signs of nitrogen deficiency even when adequate nitrogen is present in the soil.

Most plants take nitrogen from the soil continuously throughout their lives, and nitrogen demand usually increases as plant size increases. 
A plant supplied with adequate nitrogen grows rapidly and produces large amounts of succulent, green foliage. 
Providing adequate nitrogen allows an annual crop, such as corn, to grow to full maturity, rather than delaying it. 

A nitrogen-deficient plant is generally small and develops slowly because it lacks the nitrogen necessary to manufacture adequate structural and genetic materials. 

It is usually pale green or yellowish because it lacks adequate chlorophyll. 
Older leaves often become necrotic and die as the plant moves nitrogen from less important older tissues to more important younger ones.

On the other hand, some plants may grow so rapidly when supplied with excessive nitrogen that they develop protoplasm faster than they can build sufficient supporting material in cell walls. 

Such plants are often rather weak and may be prone to mechanical injury. 
Development of weak straw and lodging of small grains are an example of such an effect.

FERTILIZER MANAGEMENT
NITROGEN CYCLE:
Nitrogen fertilizer rates are determined by the crop to be grown, yield goal and quantity of nitrogen that might be provided by the soil. 
Rates needed to achieve different yields with different crops vary by region, and such decisions are usually based on local recommendations and experience.

FACTORS THAT DETERMINE THE QUANTITY OF NITROGEN SUPPLIED BY THE SOIL
The quantity of nitrogen released from the soil organic matter.

The quantity of nitrogen released by decomposition of residues of the previous crop.
Any nitrogen supplied by previous applications of organic waste.

Any nitrogen carried over from previous fertilizer applications.
Such contributions can be determined by taking nitrogen credits (expressed in lb/acre) for these variables. 

For example, corn following alfalfa usually requires less additional nitrogen than corn following corn, and less nitrogen fertilizer is needed to reach a given yield goal when manure is applied. 
As with rates, credits are usually based on local conditions.

Soil testing is being suggested more often as an alternative to taking nitrogen credits. 
Testing soils for nitrogen has been a useful practice in the drier regions of the Great Plains for many years, and in that region, fertilizer rates are often adjusted to account for NO₃⁻ found in the soil prior to planting. 

In recent years, there has been some interest in testing cornfields for NO₃⁻ in the more humid regions of the eastern United States and Canada, utilizing samples taken in late spring, after crop emergence, rather than before planting. 

This strategy, the pre-side-dress nitrogen soil test (PSNT), has received a great deal of publicity and seems to provide some indication of whether additional side-dressed nitrogen is needed or not.

FERTILIZER PLACEMENT
Placement decisions should maximize availability of nitrogen to crops and minimize potential losses. 
A plant’s roots usually will not grow across the root zone of another plant, so nitrogen must be placed where all plants have direct access to it. 

Broadcast applications accomplish this objective. 
Banding does also when all crop rows are directly next to a band. 

For corn, banding anhydrous ammonia or urea ammonium nitrate (UAN) in alternate row middles is usually as effective as banding in each middle because all rows have access to the fertilizer. 

Moist soil conditions are necessary for nutrient uptake. 
Placement below the soil surface can increase nitrogen availability under dry conditions because roots are more likely to find nitrogen in moist soil with such placement. 

Injecting side-dressed UAN may produce higher corn yields than surface application in years when dry weather follows side-dressing. 
In years when rainfall occurs shortly after application, subsurface placement is not as critical. 

Subsurface placement is normally used to control nitrogen losses. 
Anhydrous ammonia must be placed and sealed below the surface to eliminate direct volatilization losses of the gaseous ammonia. 

Volatilization from urea and UAN solutions can be controlled by incorporation or injection. 
Incorporating urea materials (mechanically or by rainfall shortly after application) is especially important in no-till situations in which volatilization is aggravated by large amounts of organic material on the soil surface. 

Placing nitrogen with phosphorus often increases phosphorus uptake, particularly when nitrogen is in the NH₄⁺ form and the crop is growing in an alkaline soil. 

The reasons for the effect are not completely clear, but may be due to nitrogen increasing root activity and potential for phosphorus uptake, and nitrification of NH₄⁺ providing acidity, which enhances phosphorus solubility. 

TIMING OF NUTRIENT APPLICATION
Timing has a major effect on the efficiency of nitrogen management systems. 
Nitrogen should be applied to avoid periods of significant loss and to provide adequate nitrogen when the crop needs Nitrogen most. 

Wheat takes up most of its nitrogen in the spring and early summer, and corn absorbs most nitrogen in midsummer, so ample availability at these times is critical. 

If losses are expected to be minimal, or can be effectively controlled, applications before or immediately after planting are effective for both crops. 

If significant losses, particularly those due to denitrification or leaching, are anticipated, split applications, in which much of the nitrogen is applied after crop emergence, can be effective in reducing losses. 

Fall applications for corn can be used on well-drained soils, particularly if the nitrogen is applied as anhydrous ammonia amended; however, fall applications should be avoided on poorly drained soils, due to an almost unavoidable potential for significant denitrification losses. 

When most of a crop’s nitrogen supply will be applied after significant crop growth or positioned away from the seed row (anhydrous ammonia or UAN banded in row middles), applying some nitrogen easily accessible to the seedling at planting ensures that the crop will not become nitrogen deficient before gaining access to the main supply of nitrogen. 

MINIMIZING FERTILIZER LOSSES
The major mechanisms for nitrogen fertilizer loss are denitrification, leaching and volatilization. 
Denitrification and leaching occur under very wet soil conditions, while volatilization is most common when soils are only moist and are drying.

PRACTICES FOR AVOIDING NITROGEN FERTILIZER LOSSES
Using an NH₄⁺ source of nitrogen acidifies the soil because the hydrogen ions (H⁺) released during nitrification of the NH₄⁺ are the major cause of acidity in soils. 
Over time, acidification and lowering of soil pH can become significant.

Nitrogen fertilizers containing NO₃⁻ but no NH₄⁺ make the soil slightly less acidic over time, but are generally used in much lesser quantities than the others. 
The acidification due to NH₄⁻ nitrogen is a significant factor in the acidification of agricultural fields, but can easily be controlled by normal liming practices.

FERTILIZING LEGUMES WITH NITROGEN
Because the Rhizobia bacteria that infect legume roots normally supply adequate nitrogen to the host plant, well-nodulated legumes rarely respond to additions of nitrogen fertilizer. 

Occasionally, however, soybeans may respond to applications of nitrogen late in the season, presumably because nitrogen fixation in the nodules has declined significantly. 
Such responses are quite erratic, though, and late-season applications of nitrogen to soybeans are not routinely recommended. 

The amount of atmospheric nitrogen fixed by non-symbiotic soil organisms varies with soil types, organic matter present and soil pH. 

NITROGEN IN PLANTS
Healthy plants often contain 3 to 4 percent nitrogen in their above-ground tissues. 
This is a much higher concentration compared to other nutrients. 

Carbon, hydrogen and oxygen, nutrients that don’t play a significant role in most soil fertility management programs, are the only other nutrients present in higher concentrations.

Nitrogen is so vital because Nitrogen is a major component of chlorophyll, the compound by which plants use sunlight energy to produce sugars from water and carbon dioxide (i.e., photosynthesis). 
Nitrogen is also a major component of amino acids, the building blocks of proteins. 

Without proteins, plants wither and die. 
Some proteins act as structural units in plant cells while others act as enzymes, making possible many of the biochemical reactions on which life is based. 

Nitrogen is a component of energy-transfer compounds, such as ATP (adenosine triphosphate). 
ATP allows cells to conserve and use the energy released in metabolism. 

Finally, nitrogen is a significant component of nucleic acids such as DNA, the genetic material that allows cells (and eventually whole plants) to grow and reproduce. 
Without nitrogen, there would be no life as we know it. 

SOIL NITROGEN
Soil nitrogen exists in three general forms: organic nitrogen compounds, ammonium (NH₄⁺) ions and nitrate (NO₃⁻) ions.

At any given time, 95 to 99 percent of the potentially available nitrogen in the soil is in organic forms, either in plant and animal residues, in the relatively stable soil organic matter, or in living soil organisms, mainly microbes such as bacteria.

This nitrogen is not directly available to plants, but some can be converted to available forms by microorganisms. 
A very small amount of organic nitrogen may exist in soluble organic compounds, such as urea, that may be slightly available to plants.

The majority of plant-available nitrogen is in the inorganic forms NH₄⁺ and NO₃⁻ (sometimes called mineral nitrogen). Ammonium ions bind to the soil's negatively charged cation exchange complex (CEC) and behave much like other cations in the soil. 

Nitrate ions do not bind to the soil solids because they carry negative charges, but exist dissolved in the soil water, or precipitated as soluble salts under dry conditions.

NATURAL SOURCES OF SOIL NITROGEN
The nitrogen in soil that might eventually be used by plants has two sources: nitrogen- containing minerals and the vast storehouse of nitrogen in the atmosphere. 
The nitrogen in soil minerals is released as the mineral decomposes. 

This process is generally quite slow, and contributes only slightly to nitrogen nutrition on most soils. 
On soils containing large quantities of NH₄⁺-rich clays (either naturally occurring or developed by fixation of NH₄⁺ added as fertilizer), however, nitrogen supplied by the mineral fraction may be significant in some years.

Atmospheric nitrogen is a major source of nitrogen in soils. 
In the atmosphere, Nitrogen exists in the very inert N₂ form and must be converted before Nitrogen becomes useful in the soil. 

The quantity of nitrogen added to the soil in this manner is directly related to thunderstorm activity, but most areas probably receive no more than 20 lb nitrogen/acre per year from this source.

Bacteria such as Rhizobia that infect (nodulate) the roots of, and receive much food energy from, legume plants can fix much more nitrogen per year (some well over 100 lb nitrogen/acre). 

When the quantity of nitrogen fixed by Rhizobia exceeds that needed by the microbes themselves, Nitrogen is released for use by the host legume plant. 

This is why well-nodulated legumes do not often respond to additions of nitrogen fertilizer. 
They are already receiving enough from the bacteria.

PRODUCTION of NITROGEN:
Nitrogen gas is an industrial gas produced by the fractional distillation of liquid air, or by mechanical means using gaseous air (pressurised reverse osmosis membrane or pressure swing adsorption). 

Nitrogen gas generators using membranes or pressure swing adsorption (PSA) are typically more cost and energy efficient than bulk delivered nitrogen. 
Commercial nitrogen is often a byproduct of air-processing for industrial concentration of oxygen for steelmaking and other purposes. 

When supplied compressed in cylinders it is often called OFN (oxygen-free nitrogen). 
Commercial-grade nitrogen already contains at most 20 ppm oxygen, and specially purified grades containing at most 2 ppm oxygen and 10 ppm argon are also available.

In a chemical laboratory, Nitrogen is prepared by treating an aqueous solution of ammonium chloride with sodium nitrite.

NH4Cl + NaNO2 → N2 + NaCl + 2 H2O

Small amounts of the impurities NO and HNO3 are also formed in this reaction. 
The impurities can be removed by passing the gas through aqueous sulfuric acid containing potassium dichromate. 
Very pure nitrogen can be prepared by the thermal decomposition of barium azide or sodium azide.

2 NaN3 → 2 Na + 3 N2

Commercial production of nitrogen is largely by fractional distillation of liquefied air. 
The boiling temperature of nitrogen is −195.8 °C (−320.4 °F), about 13 °C (−23 °F) below that of oxygen, which is therefore left behind. 

Nitrogen can also be produced on a large scale by burning carbon or hydrocarbons in air and separating the resulting carbon dioxide and water from the residual nitrogen. 

On a small scale, pure nitrogen is made by heating barium azide, Ba(N3)2. 
Various laboratory reactions that yield nitrogen include heating ammonium nitrite (NH4NO2) solutions, oxidation of ammonia by bromine water, and oxidation of ammonia by hot cupric oxide.

OCCURRENCE of NITROGEN:
Nitrogen is the most common pure element in the earth, making up 78.1% of the volume of the atmosphere (75.5% by mass), around 3.89 million gigatonnes. 

Despite this, Nitrogen is not very abundant in Earth's crust, making up somewhere around 19 parts per million of this, on par with niobium, gallium, and lithium. 
(This represents 300,000 to a million gigatonnes of nitrogen, depending on the mass of the crust.) 

The only important nitrogen minerals are nitre (potassium nitrate, saltpetre) and soda nitre (sodium nitrate, Chilean saltpetre). 
However, these have not been an important source of nitrates since the 1920s, when the industrial synthesis of ammonia and nitric acid became common.

Nitrogen compounds constantly interchange between the atmosphere and living organisms. 
Nitrogen must first be processed, or "fixed", into a plant-usable form, usually ammonia. 

Some nitrogen fixation is done by lightning strikes producing the nitrogen oxides, but most is done by diazotrophic bacteria through enzymes known as nitrogenases (although today industrial nitrogen fixation to ammonia is also significant). 

When the ammonia is taken up by plants, Nitrogen is used to synthesise proteins. 
These plants are then digested by animals who use the nitrogen compounds to synthesise their proteins and excrete nitrogen-bearing waste. 

Finally, these organisms die and decompose, undergoing bacterial and environmental oxidation and denitrification, returning free dinitrogen to the atmosphere. 

Industrial nitrogen fixation by the Haber process is mostly used as fertiliser, although excess nitrogen–bearing waste, when leached, leads to eutrophication of freshwater and the creation of marine dead zones, as nitrogen-driven bacterial growth depletes water oxygen to the point that all higher organisms die. 
Furthermore, nitrous oxide, which is produced during denitrification, attacks the atmospheric ozone layer.

Many saltwater fish manufacture large amounts of trimethylamine oxide to protect them from the high osmotic effects of their environment; conversion of this compound to dimethylamine is responsible for the early odour in unfresh saltwater fish. 

In animals, free radical nitric oxide (derived from an amino acid), serves as an important regulatory molecule for circulation.

Nitric oxide's rapid reaction with water in animals results in the production of its metabolite nitrite. 
Animal metabolism of nitrogen in proteins, in general, results in the excretion of urea, while animal metabolism of nucleic acids results in the excretion of urea and uric acid. 

The characteristic odour of animal flesh decay is caused by the creation of long-chain, nitrogen-containing amines, such as putrescine and cadaverine, which are breakdown products of the amino acids ornithine and lysine, respectively, in decaying proteins.

OCCURRENCE AND DISTRIBUTION of NITROGEN:
Among the elements, nitrogen ranks sixth in cosmic abundance. 
The atmosphere of Earth consists of 75.51 percent by weight (or 78.09 percent by volume) of nitrogen; this is the principal source of nitrogen for commerce and industry. 

The atmosphere also contains varying small amounts of ammonia and ammonium salts, as well as nitrogen oxides and nitric acid (the latter substances being formed in electrical storms and in the internal combustion engine). 

Free nitrogen is found in many meteorites; in gases of volcanoes, mines, and some mineral springs; in the Sun; and in some stars and nebulae.

Nitrogen also occurs in mineral deposits of nitre or saltpetre (potassium nitrate, KNO3) and Chile saltpetre (sodium nitrate, NaNO3), but these deposits exist in quantities that are wholly inadequate for human needs. 
Another material rich in nitrogen is guano, found in bat caves and in dry places frequented by birds. 

In combination, nitrogen is found in the rain and soil as ammonia and ammonium salts and in seawater as ammonium (NH4+), nitrite (NO2−), and nitrate (NO3−) ions.

Nitrogen constitutes on the average about 16 percent by weight of the complex organic compounds known as proteins, present in all living organisms. 

The natural abundance of nitrogen in Earth’s crust is 0.3 part per 1,000. 
The cosmic abundance—the estimated total abundance in the universe—is between three and seven atoms per atom of silicon, which is taken as the standard.

India, Russia, the United States, Trinidad and Tobago, and Ukraine were the top five producers of nitrogen (in the form of ammonia) in the early 21st century.

ALLOTROPES of NITROGEN:
Atomic nitrogen, also known as active nitrogen, is highly reactive, being a triradical with three unpaired electrons. 

Free nitrogen atoms easily react with most elements to form nitrides, and even when two free nitrogen atoms collide to produce an excited N2 molecule, they may release so much energy on collision with even such stable molecules as carbon dioxide and water to cause homolytic fission into radicals such as CO and O or OH and H. 

Atomic nitrogen is prepared by passing an electric discharge through nitrogen gas at 0.1–2 mmHg, which produces atomic nitrogen along with a peach-yellow emission that fades slowly as an afterglow for several minutes even after the discharge terminates.

Given the great reactivity of atomic nitrogen, elemental nitrogen usually occurs as molecular N2, dinitrogen. 
Nitrogen is a colourless, odourless, and tasteless diamagnetic gas at standard conditions: Nitrogen melts at −210 °C and boils at −196 °C. 

Dinitrogen is mostly unreactive at room temperature, but Nitrogen will nevertheless react with lithium metal and some transition metal complexes. 
This is due to its bonding, which is unique among the diatomic elements at standard conditions in that it has an N≡N triple bond. 

Triple bonds have short bond lengths (in this case, 109.76 pm) and high dissociation energies (in this case, 945.41 kJ/mol), and are thus very strong, explaining dinitrogen's low level of chemical reactivity.

Other nitrogen oligomers and polymers may be possible. 
If they could be synthesised, they may have potential applications as materials with a very high energy density, that could be used as powerful propellants or explosives. 

Under extremely high pressures (1.1 million atm) and high temperatures (2000 K), as produced in a diamond anvil cell, nitrogen polymerises into the single-bonded cubic gauche crystal structure. 

This structure is similar to that of diamond, and both have extremely strong covalent bonds, resulting in its nickname "nitrogen diamond".

At atmospheric pressure, molecular nitrogen condenses (liquefies) at 77 K (−195.79 °C) and freezes at 63 K (−210.01 °C) into the beta hexagonal close-packed crystal allotropic form. 
Below 35.4 K (−237.6 °C) nitrogen assumes the cubic crystal allotropic form (called the alpha phase). 

Liquid nitrogen, a colourless fluid resembling water in appearance, but with 80.8% of the density (the density of liquid nitrogen at its boiling point is 0.808 g/mL), is a common cryogen. 

Solid nitrogen has many crystalline modifications. 
Nitrogen forms a significant dynamic surface coverage on Pluto and outer moons of the Solar System such as Triton. 

Even at the low temperatures of solid nitrogen it is fairly volatile and can sublime to form an atmosphere, or condense back into nitrogen frost. 
Nitrogen is very weak and flows in the form of glaciers and on Triton geysers of nitrogen gas come from the polar ice cap region.

DINITROGEN COMPLEXES:
The first example of a dinitrogen complex to be discovered was [Ru(NH3)5(N2)]2+ (see figure at right), and soon many other such complexes were discovered. 

These complexes, in which a nitrogen molecule donates at least one lone pair of electrons to a central metal cation, illustrate how N2 might bind to the metal(s) in nitrogenase and the catalyst for the Haber process: these processes involving dinitrogen activation are vitally important in biology and in the production of fertilisers.

Dinitrogen is able to coordinate to metals in five different ways. 
The more well-characterised ways are the end-on M←N≡N (η1) and M←N≡N→M (μ, bis-η1), in which the lone pairs on the nitrogen atoms are donated to the metal cation. 

The less well-characterised ways involve dinitrogen donating electron pairs from the triple bond, either as a bridging ligand to two metal cations (μ, bis-η2) or to just one (η2). 

The fifth and unique method involves triple-coordination as a bridging ligand, donating all three electron pairs from the triple bond (μ3-N2). 
A few complexes feature multiple N2 ligands and some feature N2 bonded in multiple ways. 

Since N2 is isoelectronic with carbon monoxide (CO) and acetylene (C2H2), the bonding in dinitrogen complexes is closely allied to that in carbonyl compounds, although N2 is a weaker σ-donor and π-acceptor than CO. 

Theoretical studies show that σ donation is a more important factor allowing the formation of the M–N bond than π back-donation, which mostly only weakens the N–N bond, and end-on (η1) donation is more readily accomplished than side-on (η2) donation.

Today, dinitrogen complexes are known for almost all the transition metals, accounting for several hundred compounds. 
They are normally prepared by three methods:

Replacing labile ligands such as H2O, H−, or CO directly by nitrogen: these are often reversible reactions that proceed at mild conditions.

Reducing metal complexes in the presence of a suitable coligand in excess under nitrogen gas. 
A common choice include replacing chloride ligands by dimethylphenylphosphine (PMe2Ph) to make up for the smaller number of nitrogen ligands attached than the original chlorine ligands.

Converting a ligand with N–N bonds, such as hydrazine or azide, directly into a dinitrogen ligand.
Occasionally the N≡N bond may be formed directly within a metal complex, for example by directly reacting coordinated ammonia (NH3) with nitrous acid (HNO2), but this is not generally applicable. 

Most dinitrogen complexes have colours within the range white-yellow-orange-red-brown; a few exceptions are known, such as the blue [{Ti(η5-C5H5)2}2-(N2)].

NITROGEN AND WATER:
Nutrients, such as nitrogen and phosphorus, are essential for plant and animal growth and nourishment, but the overabundance of certain nutrients in water can cause a number of adverse health and ecological effects. 

Nitrogen, in the forms of nitrate, nitrite, or ammonium, is a nutrient needed for plant growth. 
About 78% of the air that we breathe is composed of nitrogen gas, and in some areas of the United States, particularly the northeast, certain forms of nitrogen are commonly deposited in acid rain.

Of course, nitrogen is used in agriculture to grow crops, and on many farms the landscape has been greatly modified to maximize farming output. 
Fields have been leveled and modified to efficiently drain off excess water that may fall as precipitation or from irrigation practices.

Nitrogen is transported in water and can also be stored in sediments and plants, including algae. 
Algal blooms (both macroalgae and microalgae) and increased plant growth occur in areas with increased nutrients and high light availability.

Different parts of the landscape process dissolved and particulate nitrogen differently depending on the soils, vegetation, hydrology, and sources of nitrogen and carbon. 
Some land units and land uses contribute nitrogen to waterways, others, like wetlands, remove it.

Nitrogen, total(organic and inorganic) is not a single chemical substance but a wide range of compounds of nitrogen which can be used directly by plants as a source of nitrogen required for their nutrition directly, or else may be converted into forms that plants can use in the environment.

Nitrogen gas makes up about 80% of the atmosphere and is fairly non-reactive. 
Before living organisms can make use of nitrogen as a nutrient element, Nitrogen has to be in a soluble form, such as nitrate or ammonia (inorganic forms) or in organic forms such as proteins which are readily converted into inorganic forms. 

The key inorganic form of nitrogen is ammonia, a highly reactive gas in its pure form, but which is usually found as ammonium ions when dissolved in water. 

Ammonium ions are the form of nitrogen used by plants and animals to make amino acids which are then built up into the structural proteins and enzymes that are the key components of all living organisms. 

The other main inorganic form of nitrogen is the nitrate ion (derived from nitric acid). 
Nitrate and ammonia are inter-converted in the water and soil by micro-organisms. 
These inorganic forms are taken up by plants and converted to proteins. 

Higher animals could not use inorganic forms of nitrogen for their nutrition but instead depend on organic forms already made by the plants and other animals that they eat. 

Proteins in the animal diet are first broken down into their constituent amino acids and are then reassembled into the required animal proteins and other cell constituents. 

Waste organic nitrogen compounds in animals are converted into ammonia, and then excreted in the urine as urea and other simple organic forms. 
These substances then enter the environment where they are converted into nitrate, ammonium ions and nitrogen gas. 

Some specialised bacteria can also 'fix' nitrogen gas from the atmosphere into ammonium ions. 
The above process is called the nitrogen cycle. 
In addition to ammonium and nitrate ions, traces of unstable inorganic nitrogen compounds such as nitrite ions may occasionally be found.

SOURCES of NITROGEN:
Although nitrogen is abundant naturally in the environment, Nitrogen is also introduced through sewage and fertilizers. 

Chemical fertilizers or animal manure is commonly applied to crops to add nutrients. 
It may be difficult or expensive to retain on site all nitrogen brought on to farms for feed or fertilizer and generated by animal manure. 

Unless specialized structures have been built on the farms, heavy rains can generate runoff containing these materials into nearby streams and lakes. 
Wastewater-treatment facilities that do not specifically remove nitrogen can also lead to excess levels of nitrogen in surface or groundwater.

Nitrate can get into water directly as the result of runoff of fertilizers containing nitrate. 
Some nitrate enters water from the atmosphere, which carries nitrogen-containing compounds derived from automobiles and other sources. 

More than 3 million tons of nitrogen are deposited in the United States each year from the atmosphere, derived either naturally from chemical reactions or from the combustion of fossil fuels, such as coal and gasoline. 

Nitrate can also be formed in water bodies through the oxidation of other forms of nitrogen, including nitrite, ammonia, and organic nitrogen compounds such as amino acids. 
Ammonia and organic nitrogen can enter water through sewage effluent and runoff from land where manure has been applied or stored.

Nitrogen appears as a colorless odorless gas. Noncombustible and nontoxic. 
Makes up the major portion of the atmosphere, but will not support life by itself. 
Used in food processing, in purging air conditioning and refrigeration systems, and in pressurizing aircraft tires. 

NITRIDES, AZIDES, and NITRIDO COMPLEXES of NITROGEN:
Nitrogen bonds to almost all the elements in the periodic table except the first three noble gases, helium, neon, and argon, and some of the very short-lived elements after bismuth, creating an immense variety of binary compounds with varying properties and applications. 

Many binary compounds are known: with the exception of the nitrogen hydrides, oxides, and fluorides, these are typically called nitrides. 
Many stoichiometric phases are usually present for most elements (e.g. MnN, Mn6N5, Mn3N2, Mn2N, Mn4N, and MnxN for 
9.2 < x < 25.3). 

They may be classified as "salt-like" (mostly ionic), covalent, "diamond-like", and metallic (or interstitial), although this classification has limitations generally stemming from the continuity of bonding types instead of the discrete and separate types that it implies. 

They are normally prepared by directly reacting a metal with nitrogen or ammonia (sometimes after heating), or by thermal decomposition of metal amides:

3 Ca + N2 → Ca3N2
3 Mg + 2 NH3 → Mg3N2 + 3 H2 (at 900 °C)
3 Zn(NH2)2 → Zn3N2 + 4 NH3

Many variants on these processes are possible. 
The most ionic of these nitrides are those of the alkali metals and alkaline earth metals, Li3N (Na, K, Rb, and Cs do not form stable nitrides for steric reasons) and M3N2 (M = Be, Mg, Ca, Sr, Ba). 

These can formally be thought of as salts of the N3− anion, although charge separation is not actually complete even for these highly electropositive elements.
However, the alkali metal azides NaN3 and KN3, featuring the linear N−3 anion, are well-known, as are Sr(N3)2 and Ba(N3)2. 

Azides of the B-subgroup metals (those in groups 11 through 16) are much less ionic, have more complicated structures, and detonate readily when shocked.

Many covalent binary nitrides are known. 
Examples include cyanogen ((CN)2), triphosphorus pentanitride (P3N5), disulfur dinitride (S2N2), and tetrasulfur tetranitride (S4N4). 

The essentially covalent silicon nitride (Si3N4) and germanium nitride (Ge3N4) are also known: silicon nitride in particular would make a promising ceramic if not for the difficulty of working with and sintering it. 

In particular, the group 13 nitrides, most of which are promising semiconductors, are isoelectronic with graphite, diamond, and silicon carbide and have similar structures: their bonding changes from covalent to partially ionic to metallic as the group is descended. 

In particular, since the B–N unit is isoelectronic to C–C, and carbon is essentially intermediate in size between boron and nitrogen, much of organic chemistry finds an echo in boron–nitrogen chemistry, such as in borazine ("inorganic benzene"). 

Nevertheless, the analogy is not exact due to the ease of nucleophilic attack at boron due to its deficiency in electrons, which is not possible in a wholly carbon-containing ring.

The largest category of nitrides are the interstitial nitrides of formulae MN, M2N, and M4N (although variable composition is perfectly possible), where the small nitrogen atoms are positioned in the gaps in a metallic cubic or hexagonal close-packed lattice. 

They are opaque, very hard, and chemically inert, melting only at very high temperatures (generally over 2500 °C). 
They have a metallic lustre and conduct electricity as do metals. 
They hydrolyse only very slowly to give ammonia or nitrogen.

The nitride anion (N3−) is the strongest π donor known amongst ligands (the second-strongest is O2−). 
Nitrido complexes are generally made by thermal decomposition of azides or by deprotonating ammonia, and they usually involve a terminal {≡N}3− group. 

The linear azide anion (N−3), being isoelectronic with nitrous oxide, carbon dioxide, and cyanate, forms many coordination complexes. Further catenation is rare, although N4−4 (isoelectronic with carbonate and nitrate) is known.

HYDRIDES of NITROGEN:
Industrially, ammonia (NH3) is the most important compound of nitrogen and is prepared in larger amounts than any other compound, because it contributes significantly to the nutritional needs of terrestrial organisms by serving as a precursor to food and fertilisers. 

It is a colourless alkaline gas with a characteristic pungent smell. 
The presence of hydrogen bonding has very significant effects on ammonia, conferring on it its high melting (−78 °C) and boiling (−33 °C) points. 

As a liquid, it is a very good solvent with a high heat of vaporisation (enabling it to be used in vacuum flasks), that also has a low viscosity and electrical conductivity and high dielectric constant, and is less dense than water. 

However, the hydrogen bonding in NH3 is weaker than that in H2O due to the lower electronegativity of nitrogen compared to oxygen and the presence of only one lone pair in NH3 rather than two in H2O. 
It is a weak base in aqueous solution (pKb 4.74); its conjugate acid is ammonium, NH+4. 
It can also act as an extremely weak acid, losing a proton to produce the amide anion, NH−2. 

It thus undergoes self-dissociation, similar to water, to produce ammonium and amide. 
Ammonia burns in air or oxygen, though not readily, to produce nitrogen gas; it burns in fluorine with a greenish-yellow flame to give nitrogen trifluoride. 

Reactions with the other nonmetals are very complex and tend to lead to a mixture of products. 
Ammonia reacts on heating with metals to give nitrides.

Many other binary nitrogen hydrides are known, but the most important are hydrazine (N2H4) and hydrogen azide (HN3). 
Although it is not a nitrogen hydride, hydroxylamine (NH2OH) is similar in properties and structure to ammonia and hydrazine as well. 

Hydrazine is a fuming, colourless liquid that smells similarly to ammonia. 
Its physical properties are very similar to those of water (melting point 2.0 °C, boiling point 113.5 °C, density 1.00 g/cm3). 

Despite it being an endothermic compound, it is kinetically stable. 
It burns quickly and completely in air very exothermically to give nitrogen and water vapour. 
It is a very useful and versatile reducing agent and is a weaker base than ammonia. 

It is also commonly used as a rocket fuel.
Hydrazine is generally made by reaction of ammonia with alkaline sodium hypochlorite in the presence of gelatin or glue:

NH3 + OCl− → NH2Cl + OH−
NH2Cl + NH3 → N2H+5 + Cl− (slow)
N2H+5 + OH− → N2H4 + H2O (fast)

(The attacks by hydroxide and ammonia may be reversed, thus passing through the intermediate NHCl− instead.) 
The reason for adding gelatin is that it removes metal ions such as Cu2+ that catalyses the destruction of hydrazine by reaction with monochloramine (NH2Cl) to produce ammonium chloride and nitrogen.

Hydrogen azide (HN3) was first produced in 1890 by the oxidation of aqueous hydrazine by nitrous acid. 
It may be considered the conjugate acid of the azide anion, and is similarly analogous to the hydrohalic acids.

HALIDES and OXOHALIDES of NITROGEN:
All four simple nitrogen trihalides are known. 
A few mixed halides and hydrohalides are known, but are mostly unstable; examples include NClF2, NCl2F, NBrF2, NF2H, NFH2, NCl2H, and NClH2.

Five nitrogen fluorides are known. 
Nitrogen trifluoride (NF3, first prepared in 1928) is a colourless and odourless gas that is thermodynamically stable, and most readily produced by the electrolysis of molten ammonium fluoride dissolved in anhydrous hydrogen fluoride. 

Like carbon tetrafluoride, it is not at all reactive and is stable in water or dilute aqueous acids or alkalis. 
Only when heated does it act as a fluorinating agent, and it reacts with copper, arsenic, antimony, and bismuth on contact at high temperatures to give tetrafluorohydrazine (N2F4). 

The cations NF+4 and N2F+3 are also known (the latter from reacting tetrafluorohydrazine with strong fluoride-acceptors such as arsenic pentafluoride), as is ONF3, which has aroused interest due to the short N–O distance implying partial double bonding and the highly polar and long N–F bond. 

Tetrafluorohydrazine, unlike hydrazine itself, can dissociate at room temperature and above to give the radical NF2•. 
Fluorine azide (FN3) is very explosive and thermally unstable. 

Dinitrogen difluoride (N2F2) exists as thermally interconvertible cis and trans isomers, and was first found as a product of the thermal decomposition of FN3.

Nitrogen trichloride (NCl3) is a dense, volatile, and explosive liquid whose physical properties are similar to those of carbon tetrachloride, although one difference is that NCl3 is easily hydrolysed by water while CCl4 is not. 

It was first synthesised in 1811 by Pierre Louis Dulong, who lost three fingers and an eye to its explosive tendencies. 
As a dilute gas it is less dangerous and is thus used industrially to bleach and sterilise flour. 
Nitrogen tribromide (NBr3), first prepared in 1975, is a deep red, temperature-sensitive, volatile solid that is explosive even at −100 °C. 

Nitrogen triiodide (NI3) is still more unstable and was only prepared in 1990. 
Its adduct with ammonia, which was known earlier, is very shock-sensitive: it can be set off by the touch of a feather, shifting air currents, or even alpha particles. 

For this reason, small amounts of nitrogen triiodide are sometimes synthesised as a demonstration to high school chemistry students or as an act of "chemical magic". 
Chlorine azide (ClN3) and bromine azide (BrN3) are extremely sensitive and explosive.

Two series of nitrogen oxohalides are known: the nitrosyl halides (XNO) and the nitryl halides (XNO2). 
The first are very reactive gases that can be made by directly halogenating nitrous oxide. 
Nitrosyl fluoride (NOF) is colourless and a vigorous fluorinating agent. 

Nitrosyl chloride (NOCl) behaves in much the same way and has often been used as an ionising solvent. 
Nitrosyl bromide (NOBr) is red. The reactions of the nitryl halides are mostly similar: nitryl fluoride (FNO2) and nitryl chloride (ClNO2) are likewise reactive gases and vigorous halogenating agents.

OXIDES of NITROGEN:
Nitrogen forms nine molecular oxides, some of which were the first gases to be identified: N2O (nitrous oxide), NO (nitric oxide), N2O3 (dinitrogen trioxide), NO2 (nitrogen dioxide), N2O4 (dinitrogen tetroxide), N2O5 (dinitrogen pentoxide), N4O (nitrosylazide), and N(NO2)3 (trinitramide). 

All are thermally unstable towards decomposition to their elements. 
One other possible oxide that has not yet been synthesised is oxatetrazole (N4O), an aromatic ring.

Nitrous oxide (N2O), better known as laughing gas, is made by thermal decomposition of molten ammonium nitrate at 250 °C. This is a redox reaction and thus nitric oxide and nitrogen are also produced as byproducts. 

It is mostly used as a propellant and aerating agent for sprayed canned whipped cream, and was formerly commonly used as an anaesthetic. 
Despite appearances, it cannot be considered to be the anhydride of hyponitrous acid (H2N2O2) because that acid is not produced by the dissolution of nitrous oxide in water. 

It is rather unreactive (not reacting with the halogens, the alkali metals, or ozone at room temperature, although reactivity increases upon heating) and has the unsymmetrical structure N–N–O (N≡N+O−↔−N=N+=O): above 600 °C it dissociates by breaking the weaker N–O bond. 

Nitric oxide (NO) is the simplest stable molecule with an odd number of electrons. 
In mammals, including humans, it is an important cellular signaling molecule involved in many physiological and pathological processes. 

It is formed by catalytic oxidation of ammonia. 
It is a colourless paramagnetic gas that, being thermodynamically unstable, decomposes to nitrogen and oxygen gas at 1100–1200 °C. 

Its bonding is similar to that in nitrogen, but one extra electron is added to a π* antibonding orbital and thus the bond order has been reduced to approximately 2.5; hence dimerisation to O=N–N=O is unfavourable except below the boiling point (where the cis isomer is more stable) because it does not actually increase the total bond order and because the unpaired electron is delocalised across the NO molecule, granting it stability. 

There is also evidence for the asymmetric red dimer O=N–O=N when nitric oxide is condensed with polar molecules. 
It reacts with oxygen to give brown nitrogen dioxide and with halogens to give nitrosyl halides. 
It also reacts with transition metal compounds to give nitrosyl complexes, most of which are deeply coloured.

Blue dinitrogen trioxide (N2O3) is only available as a solid because it rapidly dissociates above its melting point to give nitric oxide, nitrogen dioxide (NO2), and dinitrogen tetroxide (N2O4). 

The latter two compounds are somewhat difficult to study individually because of the equilibrium between them, although sometimes dinitrogen tetroxide can react by heterolytic fission to nitrosonium and nitrate in a medium with high dielectric constant. 

Nitrogen dioxide is an acrid, corrosive brown gas. 
Both compounds may be easily prepared by decomposing a dry metal nitrate. 
Both react with water to form nitric acid. 

Dinitrogen tetroxide is very useful for the preparation of anhydrous metal nitrates and nitrato complexes, and it became the storable oxidiser of choice for many rockets in both the United States and USSR by the late 1950s. 

This is because it is a hypergolic propellant in combination with a hydrazine-based rocket fuel and can be easily stored since it is liquid at room temperature.

The thermally unstable and very reactive dinitrogen pentoxide (N2O5) is the anhydride of nitric acid, and can be made from it by dehydration with phosphorus pentoxide. 
It is of interest for the preparation of explosives. 

It is a deliquescent, colourless crystalline solid that is sensitive to light. 
In the solid state it is ionic with structure [NO2]+[NO3]−; as a gas and in solution it is molecular O2N–O–NO2. 

Hydration to nitric acid comes readily, as does analogous reaction with hydrogen peroxide giving peroxonitric acid (HOONO2). 
It is a violent oxidising agent. 
Gaseous dinitrogen pentoxide decomposes as follows:

N2O5 ⇌ NO2 + NO3 → NO2 + O2 + NO
N2O5 + NO ⇌ 3 NO2

OXOACIDS, OXOANIONS, and OXOACID SALTS of NITROGEN:
Many nitrogen oxoacids are known, though most of them are unstable as pure compounds and are known only as aqueous solution or as salts. 

Hyponitrous acid (H2N2O2) is a weak diprotic acid with the structure HON=NOH (pKa1 6.9, pKa2 11.6). 
Acidic solutions are quite stable but above pH 4 base-catalysed decomposition occurs via [HONNO]− to nitrous oxide and the hydroxide anion. 

Hyponitrites (involving the N2O2−2 anion) are stable to reducing agents and more commonly act as reducing agents themselves. 
They are an intermediate step in the oxidation of ammonia to nitrite, which occurs in the nitrogen cycle. 
Hyponitrite can act as a bridging or chelating bidentate ligand.

Nitrous acid (HNO2) is not known as a pure compound, but is a common component in gaseous equilibria and is an important aqueous reagent: its aqueous solutions may be made from acidifying cool aqueous nitrite (NO−2, bent) solutions, although already at room temperature disproportionation to nitrate and nitric oxide is significant. 
It is a weak acid with pKa 3.35 at 18 °C. 

They may be titrimetrically analysed by their oxidation to nitrate by permanganate. 
They are readily reduced to nitrous oxide and nitric oxide by sulfur dioxide, to hyponitrous acid with tin(II), and to ammonia with hydrogen sulfide. 
Salts of hydrazinium N2H+5 react with nitrous acid to produce azides which further react to give nitrous oxide and nitrogen. 

Sodium nitrite is mildly toxic in concentrations above 100 mg/kg, but small amounts are often used to cure meat and as a preservative to avoid bacterial spoilage. 
It is also used to synthesise hydroxylamine and to diazotise primary aromatic amines as follows:

ArNH2 + HNO2 → [ArNN]Cl + 2 H2O
Nitrite is also a common ligand that can coordinate in five ways. 
The most common are nitro (bonded from the nitrogen) and nitrito (bonded from an oxygen). 
Nitro-nitrito isomerism is common, where the nitrito form is usually less stable.

Nitric acid (HNO3) is by far the most important and the most stable of the nitrogen oxoacids. 
It is one of the three most used acids (the other two being sulfuric acid and hydrochloric acid) and was first discovered by the alchemists in the 13th century. 

It is made by catalytic oxidation of ammonia to nitric oxide, which is oxidised to nitrogen dioxide, and then dissolved in water to give concentrated nitric acid. 

In the United States of America, over seven million tonnes of nitric acid are produced every year, most of which is used for nitrate production for fertilisers and explosives, among other uses. 

Anhydrous nitric acid may be made by distilling concentrated nitric acid with phosphorus pentoxide at low pressure in glass apparatus in the dark. 

It can only be made in the solid state, because upon melting it spontaneously decomposes to nitrogen dioxide, and liquid nitric acid undergoes self-ionisation to a larger extent than any other covalent liquid as follows:

2 HNO3 ⇌ H2NO+3 + NO−3 ⇌ H2O + [NO2]+ + [NO3]−

Two hydrates, HNO3·H2O and HNO3·3H2O, are known that can be crystallised. 
It is a strong acid and concentrated solutions are strong oxidising agents, though gold, platinum, rhodium, and iridium are immune to attack. 

A 3:1 mixture of concentrated hydrochloric acid and nitric acid, called aqua regia, is still stronger and successfully dissolves gold and platinum, because free chlorine and nitrosyl chloride are formed and chloride anions can form strong complexes. 

In concentrated sulfuric acid, nitric acid is protonated to form nitronium, which can act as an electrophile for aromatic nitration:

HNO3 + 2 H2SO4 ⇌ NO+2 + H3O+ + 2 HSO−4

The thermal stabilities of nitrates (involving the trigonal planar NO−3 anion) depends on the basicity of the metal, and so do the products of decomposition (thermolysis), which can vary between the nitrite (for example, sodium), the oxide (potassium and lead), or even the metal itself (silver) depending on their relative stabilities. 

Nitrate is also a common ligand with many modes of coordination.
Finally, although orthonitric acid (H3NO4), which would be analogous to orthophosphoric acid, does not exist, the tetrahedral orthonitrate anion NO3−4 is known in its sodium and potassium salts:

These white crystalline salts are very sensitive to water vapour and carbon dioxide in the air:

Na3NO4 + H2O + CO2 → NaNO3 + NaOH + NaHCO3
Despite its limited chemistry, the orthonitrate anion is interesting from a structural point of view due to its regular tetrahedral shape and the short N–O bond lengths, implying significant polar character to the bonding.

ORGANIC NITROGEN COMPOUNDS:
Nitrogen is one of the most important elements in organic chemistry. 

Many organic functional groups involve a carbon–nitrogen bond, such as amides (RCONR2), amines (R3N), imines (RC(=NR)R), imides (RCO)2NR, azides (RN3), azo compounds (RN2R), cyanates and isocyanates (ROCN or RCNO), nitrates (RONO2), nitriles and isonitriles (RCN or RNC), nitrites (RONO), nitro compounds (RNO2), nitroso compounds (RNO), oximes (RCR=NOH), and pyridine derivatives. 

C–N bonds are strongly polarised towards nitrogen. 
In these compounds, nitrogen is usually trivalent (though it can be tetravalent in quaternary ammonium salts, R4N+), with a lone pair that can confer basicity on the compound by being coordinated to a proton. 

This may be offset by other factors: for example, amides are not basic because the lone pair is delocalised into a double bond (though they may act as acids at very low pH, being protonated at the oxygen), and pyrrole is not acidic because the lone pair is delocalised as part of an aromatic ring. 

The amount of nitrogen in a chemical substance can be determined by the Kjeldahl method. 
In particular, nitrogen is an essential component of nucleic acids, amino acids and thus proteins, and the energy-carrying molecule adenosine triphosphate and is thus vital to all life on Earth.

Although the other applications are important, by far the greatest bulk of elemental nitrogen is consumed in the manufacture of nitrogen compounds. 

The triple bond between atoms in the nitrogen molecules is so strong (226 kilocalories per mole, more than twice that of molecular hydrogen) that it is difficult to cause molecular nitrogen to enter into other combinations.

The chief commercial method of fixing nitrogen (incorporating elemental nitrogen into compounds) is the Haber-Bosch process for synthesizing ammonia. 

This process was developed during World War I to lessen the dependence of Germany on Chilean nitrate. 
Nitrogen involves the direct synthesis of ammonia from its elements.

With oxygen, nitrogen forms several oxides, including nitrous oxide, N2O, in which nitrogen is in the +1 oxidation state; nitric oxide, NO, in which it is in the +2 state; and nitrogen dioxide, NO2, in which it is in the +4 state. 

Sodium nitrate (NaNO3) and potassium nitrate (KNO3) are formed by the decomposition of organic matter with compounds of these metals present. 

In certain dry areas of the world these saltpeters are found in quantity and are used as fertilizers. 
Other inorganic nitrogen compounds are nitric acid (HNO3), ammonia (NH3), the oxides (NO, NO2, N2O4, N2O), cyanides (CN-), etc.

The nitrogen cycle is one of the most important processes in nature for living organisms. 
Although nitrogen gas is relatively inert, bacteria in the soil are capable of “fixing” the nitrogen into a usable form (as a fertilizer) for plants. 
In other words, Nature has provided a method to produce nitrogen for plants to grow. 

Animals eat the plant material where the nitrogen has been incorporated into their system, primarily as protein. 
The cycle is completed when other bacteria convert the waste nitrogen compounds back to nitrogen gas. 
Nitrogen is crucial to life, as it is a component of all proteins.

NITROGEN FIXATION:
Nitrogen gas (N2) makes up nearly 80% of the Earth's atmosphere, yet nitrogen is often the nutrient that limits primary production in many ecosystems. 
Why is this so? 
Because plants and animals are not able to use nitrogen gas in that form. 

For nitrogen to be available to make proteins, DNA, and other biologically important compounds, Nitrogen must first be converted into a different chemical form. 
The process of converting N2 into biologically available nitrogen is called nitrogen fixation. 

N2 gas is a very stable compound due to the strength of the triple bond between the nitrogen atoms, and Nitrogen requires a large amount of energy to break this bond. 
The whole process requires eight electrons and at least sixteen ATP molecules. 

As a result, only a select group of prokaryotes are able to carry out this energetically demanding process. 
Although most nitrogen fixation is carried out by prokaryotes, some nitrogen can be fixed abiotically by lightning or certain industrial processes, including the combustion of fossil fuels.

NITROGEN AND THE CNO CYCLE:
When the universe’s first generation of stars was born, they contained only the elements made in the big bang: hydrogen, helium, and a tiny amount of lithium.

As these stars burned, they synthesized heavier elements, such as carbon. 
Supernovae then spread the heavier elements out into galaxies where more stars were born.

Carbon from supernovae plays a crucial role in the way many second and higher generation stars burn. 
In stars whose mass is higher than about 1.1 – 1.5 times that of our sun, carbon-12 catalyzes the fusion of hydrogen to helium – i.e. carbon-12 takes part in the fusion reaction, but is not consumed by it.

As you can see on the left, carbon-12 is regenerated at the end of each cycle, the net result of which is that four hydrogen nuclei are consumed and one helium nucleus is produced. 
This reaction is called the CNO cycle.

Over time, each carbon-12 nucleus can take part in a very large number of cycles. 
A proportion of nitrogen made during the CNO cycle evades further reaction. 

At the end of a star’s life, this nitrogen may be distributed into the galaxy. 
In our solar system the nitrogen from a star that died billions of years ago ended up as an essential element in proteins and DNA and formed about 80 percent of our planet’s atmosphere.

INTERESTING FACTS ABOUT NITROGEN:
-About 2.5 percent of the weight of living organisms comes from nitrogen in organic molecules.
-Many of the molecules of life contain nitrogen. 
Nitrogen is the fourth most abundant element in the human body.

-The nitrogen compound nitroglycerin can be used for relief of angina, a life threatening heart condition.
-Neptune’s satellite Triton has five mile high, nitrogen-powered geysers.
-Nitrogen is the seventh most abundant element in the universe.

-Like Earth, Triton’s atmosphere is mainly nitrogen, but Triton is so cold the nitrogen sits on the surface as a rock-hard solid. 
The solid nitrogen allows the feeble light arriving from the sun to pass through it. 

Dark impurities in the nitrogen ice or in darker rocks below the ice warm up slightly in the sunlight, melting and vaporizing the solid nitrogen, which eventually breaks through the solid nitrogen as geysers which push ice particles one to five miles above Triton’s frozen surface.

-In 1919, the world learned for the first time that atomic nuclei could be disintegrated. 
Ernest Rutherford reported that he had bombarded nitrogen gas with alpha-particles (helium nuclei) and found hydrogen was produced. 
(Further research by Patrick Blackett showed that the alpha particles had transmuted nitrogen-14 to oxygen-17 plus hydrogen.)
-The universe’s nitrogen was made, and is being made, by the CNO cycle in stars heavier than our sun. 
 

HISTORY of NITROGEN:
Nitrogen compounds have a very long history, ammonium chloride having been known to Herodotus. 
They were well known by the Middle Ages. 

Alchemists knew nitric acid as aqua fortis (strong water), as well as other nitrogen compounds such as ammonium salts and nitrate salts. 

The mixture of nitric and hydrochloric acids was known as aqua regia (royal water), celebrated for its ability to dissolve gold, the king of metals.

The discovery of nitrogen is attributed to the Scottish physician Daniel Rutherford in 1772, who called it noxious air. Though he did not recognise it as an entirely different chemical substance, he clearly distinguished it from Joseph Black's "fixed air", or carbon dioxide. 

The fact that there was a component of air that does not support combustion was clear to Rutherford, although he was not aware that it was an element. 

Nitrogen was also studied at about the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to Nitrogen as burnt air or phlogisticated air. 

French chemist Antoine Lavoisier referred to nitrogen gas as "mephitic air" or azote, from the Greek word άζωτικός (azotikos), "no life", due to Nitrogen being asphyxiant. 
In an atmosphere of pure nitrogen, animals died and flames were extinguished. 

Though Lavoisier's name was not accepted in English since it was pointed out that all gases but oxygen are either asphyxiant or outright toxic, it is used in many languages (French, Italian, Portuguese, Polish, Russian, Albanian, Turkish, etc.; the German Stickstoff similarly refers to the same characteristic, viz. ersticken "to choke or suffocate") and still remains in English in the common names of many nitrogen compounds, such as hydrazine and compounds of the azide ion. 

Finally, it led to the name "pnictogens" for the group headed by nitrogen, from the Greek πνίγειν "to choke".
The English word nitrogen (1794) entered the language from the French nitrogène, coined in 1790 by French chemist Jean-Antoine Chaptal (1756–1832), from the French nitre (potassium nitrate, also called saltpeter) and the French suffix -gène, "producing", from the Greek -γενής (-genes, "begotten"). 

Chaptal's meaning was that nitrogen is the essential part of nitric acid, which in turn was produced from nitre. 
In earlier times, niter had been confused with Egyptian "natron" (sodium carbonate) – called νίτρον (nitron) in Greek – which, despite the name, contained no nitrate.

The earliest military, industrial, and agricultural applications of nitrogen compounds used saltpeter (sodium nitrate or potassium nitrate), most notably in gunpowder, and later as fertiliser. 

In 1910, Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", a monatomic allotrope of nitrogen. 
The "whirling cloud of brilliant yellow light" produced by his apparatus reacted with mercury to produce explosive mercury nitride.

For a long time, sources of nitrogen compounds were limited. 
Natural sources originated either from biology or deposits of nitrates produced by atmospheric reactions. 

Nitrogen fixation by industrial processes like the Frank–Caro process (1895–1899) and Haber–Bosch process (1908–1913) eased this shortage of nitrogen compounds, to the extent that half of global food production (see Applications) now relies on synthetic nitrogen fertilisers. 

At the same time, use of the Ostwald process (1902) to produce nitrates from industrial nitrogen fixation allowed the large-scale industrial production of nitrates as feedstock in the manufacture of explosives in the World Wars of the 20th century.

About four-fifths of Earth’s atmosphere is nitrogen, which was isolated and recognized as a specific substance during early investigations of the air. 

Carl Wilhelm Scheele, a Swedish chemist, showed in 1772 that air is a mixture of two gases, one of which he called “fire air,” because it supported combustion, and the other “foul air,” because it was left after the “fire air” had been used up. 
The “fire air” was, of course, oxygen and the “foul air” nitrogen. 

At about the same time, nitrogen also was recognized by a Scottish botanist, Daniel Rutherford (who was the first to publish his findings), by the British chemist Henry Cavendish, and by the British clergyman and scientist Joseph Priestley, who, with Scheele, is given credit for the discovery of oxygen. 

Later work showed the new gas to be a constituent of nitre, a common name for potassium nitrate (KNO3), and, accordingly, it was named nitrogen by the French chemist Jean-Antoine-Claude Chaptal in 1790. 

Nitrogen first was considered a chemical element by Antoine-Laurent Lavoisier, whose explanation of the role of oxygen in combustion eventually overthrew the phlogiston theory, an erroneous view of combustion that became popular in the early 18th century. 

The inability of nitrogen to support life (Greek: zoe) led Lavoisier to name it azote, still the French equivalent of nitrogen.

Nitrogen was discovered by the Scottish physician Daniel Rutherford in 1772. 
Nitrogen is the fifth most abundant element in the universe and makes up about 78% of the earth's atmosphere, which contains an estimated 4,000 trillion tons of the gas. 

Nitrogen is obtained from liquefied air through a process known as fractional distillation.
The largest use of nitrogen is for the production of ammonia (NH3). 
Large amounts of nitrogen are combined with hydrogen to produce ammonia in a method known as the Haber process. 

The French chemist Antoine Laurent Lavoisier named nitrogen azote, meaning "without life". 
The name became nitrogen, which derives from the Greek word nitron, which means "native soda" and genes, which means "forming". 

Credit for the discovery of the element is generally given to Daniel Rutherford, who found it could be separated from air in 1772.
Nitrogen was sometimes referred to as "burnt" or "dephlogisticated" air, since air that no longer contains oxygen is almost all nitrogen. 

The other gases in air are present in much lower concentrations.
Nitrogen compounds are found in foods, fertilizers, poisons, and explosives. 
Your body is 3% nitrogen by weight. 

All living organisms contain Nitrogen.
Nitrogen is responsible for the orange-red, blue-green, blue-violet, and deep violet colors of the aurora.
One way to prepare nitrogen gas is by liquefaction and fractional distillation from the atmosphere. 

Liquid nitrogen boils at 77 K (−196 °C, −321 °F). 
Nitrogen freezes at 63 K (-210.01 °C).
Liquid nitrogen is a cryogenic fluid, capable of freezing skin on contact. 

While the Leidenfrost effect protects skin from very brief exposure (less than one second), ingesting liquid nitrogen can cause severe injury. 
When liquid nitrogen is used to make ice cream, the nitrogen vaporizes. 
Nitrogen has a valence of 3 or 5. 

Nitrogen forms negatively charged ions (anions) that readily react with other nonmetals to form covalent bonds.
Saturn's largest moon, Titan, is the only moon in the solar system with a dense atmosphere. 
Nitrogen's atmosphere consists of over 98% nitrogen.

From the Latin word nitrum, Greek Nitron, native soda; and genes, forming. 
Nitrogen was discovered by chemist and physician Daniel Rutherford in 1772. 
He removed oxygen and carbon dioxide from air and showed that the residual gas would not support combustion or living organisms. 

At the same time there were other noted scientists working on the problem of nitrogen. 
These included Scheele, Cavendish, Priestley, and others. 
They called it "burnt" or" dephlogisticated air," which meant air without oxygen.

DISCOVERY of NITROGEN:
Dr. Doug Stewart
In 1674 the English physician John Mayow demonstrated that air is not a single element, it is made up of different substances. 
He did this by showing that only a part of air is combustible. 
Most of Nitrogen is not. 

Almost a century later, Scottish chemist Joseph Black carried out more detailed work on air. 
After removing oxygen and carbon dioxide, part of the air remained.

Black used burning phosphorus as the final step in oxygen removal. 
(Burning phosphorus has a very high affinity for oxygen and is efficient at removing it completely.) 
Black then assigned further study of the gases in air to his doctoral student, Daniel Rutherford. 

Rutherford built on Black’s work and in a series of steps thoroughly removed oxygen and carbon dioxide from air. 
He showed that, like carbon dioxide, the residual gas could not support combustion or living organisms. 

Unlike carbon dioxide, however, nitrogen was insoluble in water and alkali solutions. 
Rutherford reported his discovery in 1772 of ‘noxious air,’ which we now call nitrogen. 

Swedish pharmacist Carl Scheele discovered nitrogen independently, calling it spent air.
Scheele absorbed oxygen in a number of ways, including using a mixture of sulfur and iron filings and burning phosphorus. 

After removing the oxygen, he reported a residual gas which would not support combustion and had between two-thirds and three-quarters of the volume of the original air. 
Scheele published his results in 1777, although it is thought the work was carried out in 1772. 

Although Rutherford and Scheele are now jointly credited with nitrogen’s discovery, Nitrogen appears to have been discovered earlier by Henry Cavendish, but not published.

Prior to 1772 (the precise date is unknown – Priestley refers to it in his work “Experiments and Observations Made in and Before the Year 1772”) Cavendish wrote to Joseph Priestley describing ‘burnt air’.

The ‘burnt air’ had been prepared by passing air repeatedly over red hot charcoal (removing the oxygen) and then bubbling the remaining gas through a solution of caustic potash (potassium hydroxide) which would have removed the carbon dioxide.

Cavendish wrote: “The specific gravity of this air was found to differ very little from that of common air; of the two, it seemed rather lighter. 

Nitrogen extinguished flame, and rendered common air unfit for making bodies burn in the same manner as fixed air, but in a less degree, as a candle which burnt about 80″ in pure common air, and which went out immediately in common air mixed with 6/55 of fixed air, burnt about 26″ in common air mixed with the same portion of this burnt air.” 

In 1790 the French chemist Jean-Antoine-Claude Chaptal named the element ‘nitrogen’ after experiments showed it to be a constituent of nitre, as potassium nitrate was called then.

PHYSICAL and CHEMICAL PROPERTIES of NITROGEN:
Appearance: colorless gas, liquid or solid
Standard atomic weight Ar°(N): [14.00643, 14.00728] 14.007±0.001 (abridged)
Atomic number (Z): 7
Group    group: 15 (pnictogens)
Period    period: 2
Block: p-block
Electron configuration: [He] 2s2 2p3
Electrons per shell: 2, 5
Phase at STP: gas

Melting point: (N2) 63.23[2] K ​(−209.86[2] °C, ​−345.75[2] °F)
Boiling point: (N2) 77.355 K ​(−195.795 °C, ​−320.431 °F)
Density (at STP): 1.2506 g/L[3] at 0 °C, 1013 mbar
when liquid (at b.p.): 0.808 g/cm3
Triple point: 63.151 K, ​12.52 kPa
Critical point: 126.21 K, 3.39 MPa
Heat of fusion: (N2) 0.72 kJ/mol
Heat of vaporisation: (N2) 5.57 kJ/mol
Molar heat capacity: (N2) 29.124 J/(mol·K)

Oxidation states: −3, −2, −1, 0,[4] +1, +2, +3, +4, +5 (a strongly acidic oxide)
Electronegativity: Pauling scale: 3.04
Ionisation energies:
1st: 1402.3 kJ/mol
2nd: 2856 kJ/mol
3rd: 4578.1 kJ/mol
Covalent radius: 71±1 pm
Van der Waals radius: 155 pm
Natural occurrence: primordial

Crystal structure: ​hexagonalHexagonal crystal structure for nitrogen
Speed of sound: 353 m/s (gas, at 27 °C)
Thermal conductivity: 25.83×10−3 W/(m⋅K)
Magnetic ordering: diamagnetic
Molecular Weight: 28.014    
XLogP3-AA: 0.1    
Hydrogen Bond Donor Count: 0    
Hydrogen Bond Acceptor Count: 2    
Rotatable Bond Count: 0    

Exact Mass: 28.006148008    
Monoisotopic Mass: 28.006148008    
Topological Polar Surface Area: 47.6 Ų    
Heavy Atom Count: 2    
Formal Charge: 0    
Complexity: 8    
Isotope Atom Count: 0    
Defined Atom Stereocenter Count: 0    
Undefined Atom Stereocenter Count: 0    

Defined Bond Stereocenter Count: 0    
Undefined Bond Stereocenter Count: 0    
Covalently-Bonded Unit Count: 1    
Compound Is Canonicalized: Yes
Atomic number: 7
Atomic mass: 14.0067 g.mol -1
Electronegativity according to Pauling: 3.0
Density: 1.25*10-3 g.cm-3 at 20°C
Melting point: -210 °C

Boiling point: -195.8 °C
Vanderwaals radius: 0.092 nm
Ionic radius: 0.171 nm (-3) ; 0.011 (+5) ; 0.016 (+3)
Isotopes: 4
Electronic shell: [ He ] 2s22p3
Energy of first ionisation: 1402 kJ.mol -1
Energy of second ionisation: 2856 kJ.mol -1
Energy of third ionisation: 4578 kJ.mol -1
Discovered by: Rutherford in 1772

Chemical formula : N 2
Purity level : ≥ 99.9%
Relative density (air = 1) : 0.97
Aspect : colorless gas
Odor : odorless gas
Limit of flammability in air : not flammable
Appearance Form: Compressed gas
Color: colorless
Odo:r odorless
Odor Threshold: No data available

pH: No data available
Melting point/freezing point: -209,99 °C
Initial boiling point and boiling range: -195,79 °C
Flash point: Not applicable
Evaporation rate: No data available
Flammability (solid, gas): No data available
Upper/lower flammability or explosive limits: No data available
Vapor pressure: No data available
Vapor density: No data available

Relative density: 0,97 g/cm3
Water solubility: No data available
Partition coefficient: n-octanol/water: No data available
Autoignition temperature: No data available
Decomposition temperature: No data available
Viscosity: No data available
Explosive properties: No data available
Oxidizing properties: No data available
Other safety information: No data available

Atomic number (number of protons in the nucleus): 7
Atomic symbol (on the Periodic Table of Elements): N
Atomic weight (average mass of the atom): 14.0067
Density: 0.0012506 grams per cubic centimeter
Phase at room temperature: Gas
Melting point: minus 321 degrees Fahrenheit (minus 210 degrees Celsius)
Boiling point: minus 320.42 F (minus 195.79 C)
Number of isotopes (atoms of the same element with a different number of neutrons): 16 including 2 stable ones
Most common isotopes: Nitrogen-14 (Abundance: 99.63 percent)

FIRST AID MEASURES of NITROGEN:
-Description of first-aid measures:
*General advice:
Consult a physician. 

*If inhaled:
If breathed in, move person into fresh air. 
Consult a physician.

*In case of skin contact:
Wash off with soap and plenty of water. 
Consult a physician.

*In case of eye contact:
Flush eyes with water as a precaution.

*If swallowed:
Rinse mouth with water. 
Consult a physician.

-Indication of any immediate medical attention and special treatment needed:
No data available

ACCIDENTAL RELEASE MEASURES of NITROGEN:
-Environmental precautions:
Do not let product enter drains.

-Methods and materials for containment and cleaning up:
Clean up promptly by sweeping or vacuum.

FIRE FIGHTING MEASURES of NITROGEN:
-Extinguishing media:
*Suitable extinguishing media:
Use water spray, alcohol-resistant foam, dry chemical or carbon dioxide.

-Further information:
Use water spray to cool unopened containers.

EXPOSURE CONTROLS/PERSONAL PROTECTION of NITROGEN:
-Control parameters:
--Ingredients with workplace control parameters:
-Exposure controls:
--Appropriate engineering controls:
Handle in accordance with good industrial hygiene and safety practice. 
Wash hands before breaks and at the end of workday.

--Personal protective equipment:
*Eye/face protection:
Use equipment for eye protection.

*Skin protection:
Handle with gloves. 
Wash and dry hands.

-Control of environmental exposure:
Do not let product enter drains.

HANDLING and STORAGE of NITROGEN:
-Conditions for safe storage, including any incompatibilities:
Store in cool place. 
Keep container tightly closed in a dry and well-ventilated place.
Contents under pressure.

STABILITY and REACTIVITY of NITROGEN:
-Reactivity:
No data available

-Chemical stability:
Stable under recommended storage conditions.

-Possibility of hazardous reactions:
No data available

-Conditions to avoid:
No data available

SYNONYMS:
N
Nitrogen
Nitrogen gas
Molecular nitrogen
Dinitrogen
Nitrogen-14
nitrogen molecule
Diatomic nitrogen
Nitrogen, liquid
UNII-N762921K75
N2
CHEBI:17997
N762921K75
MOL Nitrogen
Nitrogeno
Nitrogen [NF]
Nitrogen (liquified)
HSDB 5060
UN1066
UN1977
Nitrogen (TN)
Nitrogen, elemental
Nitrogen (n2)
Nitrogen, compressed
Nitrogen (JP17/NF)
Nitrogen, >=99.998%
Nitrogen, >=99.999%
CHEMBL142438
INS NO.941
DTXSID4036304
INS-941
DB09152
UN 1066
UN 1977
E941
Nitrogen, Messer(R) CANGas, 99.999%
E-941
C00697
D00083
Nitrogen, refrigerated liquid (cryogenic liquid)
Nitrogen, compressed [UN1066] [Nonflammable gas]
Q2370426
UNII-K21NZZ5Y0B component IJGRMHOSHXDMSA-UHFFFAOYSA-N
Nitrogen, refrigerated liquid (cryogenic liquid) [UN1977] [Nonflammable gas]
N#N