Nitrogen
Nitrogen is a chemical element with atomic number 7, symbol N, its atomic mass is 14.0067, and under conditions normal forms a diatomic gas (diatomic or molecular nitrogen) that constitutes about 78% of atmospheric air. Formerly it was called azote (symbol Az).
Nitrogen is the lightest member of group 15 of the periodic table, often called pnicogen. It is a common element in the universe, estimated to be about seventh in total abundance in the Milky Way and Solar System. At standard temperature and pressure, two atoms of the element unite to form dinitrogen, a colorless, odorless gas with the formula N2. Dinitrogen forms about 78% of Earth's atmosphere, making it the most abundant uncombined element. Nitrogen is present in all organisms, mainly in amino acids (and therefore in proteins), in 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, to the biosphere and organic compounds, and then back to the atmosphere.
Many industrially important compounds, such as ammonia, nitric acid, organic nitrates (propellants and explosives), and cyanide, contain nitrogen. The very strong triple bond of elemental nitrogen (N≡N), the second strongest bond in any diatomic molecule after carbon monoxide (CO), dominates nitrogen chemistry. This causes difficulties for both organisms and industry in converting N2 into useful compounds, but at the same time means that burning, exploding or decomposing nitrogen compounds to forming nitrogen gas release large amounts of often useful energy. Synthetically produced ammonia and nitrates are key fertilizers for the industry, and nitrates from fertilizers are key pollutants in the eutrophication of aquatic systems.
Apart from its use in fertilizers and energy stores, nitrogen is a constituent of organic compounds as diverse as the Kevlar used in high-strength fabrics and the cyanoacrylate used in superglue. Nitrogen is a constituent of all major drug classes, including antibiotics. Many drugs are mimics or prodrugs of natural nitrogen-containing cell signaling molecules: for example, the organic nitrates nitroglycerin and nitroprusside control blood pressure by being metabolized to nitric oxide. Many notable nitrogen-containing drugs, such as natural caffeine and synthetic morphine or amphetamines, act on receptors for animal neurotransmitters.
History
Nitrogen compounds have a very long history, since ammonium chloride was known to Herodotus. They were already well known in the Middle Ages. Alchemists knew nitric acid as aqua fortis (strong water), as well as other nitrogenous compounds such as ammonium salts and nitrate salts. The mixture of nitric and hydrochloric acid was known as aqua regia (royal water), famous 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. Although he did not recognize it as an entirely different chemical substance, he clearly distinguished it from "fixed air", or carbon dioxide, by Joseph Black. The fact that there was a component of air that did not support combustion was clear to Rutherford, although he was not aware that it was an element. Nitrogen was also studied around the same time by Carl Wilhelm Scheele, Henry Cavendish, and Joseph Priestley, who referred to it as the burnt air or phlogiston theory. French chemist Antoine Lavoisier referred to nitrogen gas as "mephitic air" or scourge, from the word άζωτικός from Greek. (azotikos), "lifeless", because it is mostly inert. In an atmosphere of pure nitrogen, the animals died and the flames were extinguished. Although Lavoisier's name was not accepted in English, since it was pointed out that almost all gases (in fact, with the sole exception of oxygen) are mephitic, it is used in many languages (French, Italian, Portuguese, Polish, Russian, Albanian, Turkish, etc. The German Stickstoff also refers to the same feature, i.e. ersticken "to choke or suffocate") and still remains in English in the common names of many nitrogen compounds, such as hydrazine and azide ion compounds. Eventually, it gave rise to the name "pnicogens" 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 the French chemist Jean-Antoine Chaptal (1756-1832), from the French nitre (potassium nitrate, also called saltpeter) and the French suffix -gène, "produce", from the Greek -γενής (-genes, "begotten& #3. 4;). Chaptal's meaning was that nitrogen is the essential part of nitric acid, which in turn was produced from potassium nitrate. In earlier times, nitro had been confused with "natron" Egyptian (sodium carbonate) - called νίτρον (nitron) in Greek - which, despite the name, did not contain nitrate.
The first military, industrial, and agricultural applications of nitrogen compounds used saltpeter (sodium nitrate or potassium nitrate), primarily in gunpowder, and later as a fertilizer. In 1910 Lord Rayleigh discovered that an electrical discharge in nitrogen gas produced "active nitrogen", a monatomic allotrope of nitrogen. produced by the apparatus he reacted with mercury to produce explosive mercury nitride.
For a long time, sources of nitrogen compounds were limited. The natural sources came from biology or from nitrate deposits produced by atmospheric reactions. Nitrogen fixation through industrial processes such as the Frank-Caro process (1895-1899) and the Haber-Bosch process (1908-1913) alleviated this shortage of nitrogenous compounds, to the point that half of the world's food production depends on now from synthetic nitrogenous fertilizers. At the same time, the use of Ostwald's method (1902) to produce nitrates from industrial nitrogen fixation enabled large-scale industrial production of nitrates as a raw material in the manufacture of explosives in World War of the 20th century.
Properties
Atomics
A nitrogen atom has seven electrons. In the basic state, they are arranged in the electronic configuration 1s2
2s2
2p1
x2p1
y2p1
z. Therefore, it has five valence electrons in the 2s and 2p orbitals, three of which (the p electrons) are unpaired. It has one of the highest electronegativities among the elements (3.04 on the Pauling scale), second only to 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 are in fact on the Allen scale.) Following periodic trends, its covalent radius of 71 pm is smaller than that of boron (84 pm) and carbon (76 pm), while it is greater 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). The first three ionization 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 large numbers, nitrogen does not have a simple cationic chemistry.
The lack of radial nodes in the 2p subsphere 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 radius very similar to the 2s shell, facilitating orbital hybridization. It also gives rise to very large attractive electrostatic forces between the nucleus and the valence electrons of the 2s and 2p shells, giving rise to very high electronegativities. Hypervalency is almost unheard of in 2p elements for the same reason, as the high electronegativity makes it difficult for a small nitrogen atom to be a central atom in a three-center, four-electron bond, as it would tend to pull electrons strongly toward it. itself. Thus, despite nitrogen's position at the head of group 15 in the periodic table, its chemistry shows vast differences from that of its heavier congeners phosphorus, arsenic, antimony, and bismuth.
Isotopes
Nitrogen has two stable isotopes: 14N and 15N. The former is much more common, making up 99.634% of natural nitrogen, and the latter (which is slightly heavier) makes up the remaining 0.366%. This leads to an atomic weight of about 14.007 u. Both stable isotopes occur in the CNO cycle in stars, but 14N is more common since its capture of neutrons is the step that limit speed. 14N is one of the five stable odd nuclides. (a nuclide that has an odd number of protons and neutrons); the other four are 2H, 6Li, 10B, and 180mTa.
The relative abundance of 14N and 15N is nearly 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 (for example, assimilation, nitrification, and denitrification) strongly control nitrogen dynamics in the soil. These reactions typically result in 15N enrichment of the growing medium and product depletion.
The heavy isotope 15N was first discovered by S. M. Naudé in 1929, shortly after the discovery of the heavy isotopes of the neighboring elements oxygen and carbon. Shows one of the cross sections lowest thermal neutron capture of all isotopes. It is frequently used in nuclear magnetic resonance (NMR) to determine the structures of nitrogen-containing molecules, due to its half-fractional nuclear spin, which offers advantages for NMR, such as a narrower linewidth. 14N, while also theoretically usable, has an integer nuclear spin of one and therefore has a quadrupole moment which leads to broader and less useful spectra. sup>15N does, however, have complications not found on NMR spectroscopy of the more common 1H and 13C. The low natural abundance of 15N (0.36%) significantly reduces sensitivity, a problem that is compounded by its low gyromagnetic ratio, (only 10.14% of that of 1H). As a result, the signal-to-noise ratio for 1H is about 300 times that of 15N at the same magnetic field strength. This can be alleviated to some extent by isotopic enrichment of 15N by chemical exchange or fractional distillation. Compounds enriched with 15N have the advantage that, under standard conditions, they do not undergo chemical exchange of their nitrogen atoms with atmospheric nitrogen, unlike compounds with isotopes of hydrogen, carbon and marked oxygen 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.
Application
The most important commercial application of diatomic nitrogen is to obtain ammonia by the Haber process. The ammonia is later used in the manufacture of fertilizers and nitric acid.
Salts of nitric acid include important compounds such as potassium nitrate (nitro or saltpeter used in the manufacture of gunpowder) and ammonium nitrate fertilizer.
Organic nitrogen compounds such as nitroglycerin and trinitrotoluene are often explosive. Hydrazine and its derivatives are used as rocket fuel.
The cycle of this element is much more complex than that of carbon, since it is present in the atmosphere not only as N2 (78%) but also in a wide variety of compounds. It can be found mainly as N2O, NO and NO2, the so-called NOx. It also forms other combinations with oxygen such as N2O3 and N2O5 (anhydrides), "precursors" of nitrous and nitric acids. With hydrogen it forms ammonia (NH3), a gaseous compound under normal conditions.
Being a non-reactive gas, nitrogen is used industrially to create protective atmospheres and as a cryogenic gas to obtain temperatures of the order of 78 K easily and cheaply. It is even used to inflate the tires on aircraft landing gear, avoiding water condensation at high altitudes or its combustion when landing.
Etymology
Nitrogen (from the Latin nitrum -i, in turn from the Greek νίτρον, "nitro" -a name that has historically been used loosely to refer to various nitrogen-containing sodium and potassium compounds, and -geno, from the Greek root γεν-, "generate&# 34;; that is, "that generates saltpeter") was formally discovered by Daniel Rutherford in 1772, when making known some of its properties (he called it "phlogisticated air", based on from what he observed in his experiment that year). However, around the same time, Carl Wilhelm Scheele, who isolated him, Henry Cavendish and Joseph Priestley, also dedicated themselves to the study of it.
Nitrogen is such an inert gas that Antoine Lavoisier referred to it by the name scourge, from the Greek azoe, meaning "lifeless" (or maybe he called it that because he wasn't fit to breathe). It was classified among the permanent gases, especially since Michael Faraday failed to see it as a liquid at 50 atmospheres (atm) and –110 °C until the experiments of Raoul Pictet and Louis Paul Cailletet, who in 1877 succeeded in liquefy it.
Nitrogen compounds were already known in the Middle Ages; thus, alchemists called nitric acid aqua fortis and aqua regia (aqua regia) the mixture of nitric acid and hydrochloric acid, a mixture known for its ability to dissolve gold and the platinum.
Abundance and obtaining
Nitrogen is the main component of the earth's atmosphere (78.1% by volume) and is obtained for industrial uses from the distillation of liquid air. It is also present in animal remains, for example guano, usually in the form of urea, uric acid, and compounds of both. Due to deficiency, it causes lack of muscle relaxation, problems in the cardiovascular system, in the central and peripheral nervous systems.
It also occupies 3% of the elemental composition of the human body.
Nitrogen-containing compounds have been observed in outer space, and the isotope Nitrogen-14 is created in the nuclear fusion processes of stars.
Most of the nitrogen used in the chemical industry is obtained by fractional distillation of liquid air, and is used to synthesize ammonia. A wide variety of chemical products are prepared from this ammonia.
Compounds
With hydrogen it forms ammonia (NH3), nitrites (NO2), nitrates (NO3), nitric acids (HNO3), hydrazine (N2H4) and hydrogen azide (N3 >H, also known as hydrogen azide or hydrazoic acid). Liquid ammonia, amphoteric like water, acts like a base in an aqueous solution, forming ammonium ions (NH4+), and behaves like an acid in the absence of water, donating a proton to a base and giving rise to the amide anion (NH2). Long chain and cyclic nitrogen compounds are known, but they are highly unstable.
With the halogens it forms: NF3, NF2Cl, NFCl2, NCl3, NBr3.6 NH3, NI3.6 NH3, N2 >F4, N2F2 (cis and trans), N3F, N3Cl, N 3Bry N3I.
With oxygen it forms several oxides that we have already named: nitrous or laughing gas, nitric and nitrogen dioxide. They are the product of combustion processes contributing to the appearance of polluting episodes of photochemical smog. Other oxides are dinitrogen trioxide (N2O3) and dinitrogen pentoxide (N2O5 >), both very unstable and explosive.
Biological importance
Nitrogen is an essential component of amino acids and nucleic acids vital to living things. Of all the mineral nutrients, it is the one that has the greatest effect on plant growth and, therefore, on the primary productivity of ecosystems, which in turn affects all the organisms that depend on them: The increase in crop yields since nitrogen fertilizers began to be used in the 19th century demonstrates this. Despite the large amount of atmospheric nitrogen, this element is limiting: few organisms can assimilate it in this form. Plants can only assimilate it efficiently in the form of ammonium ions (NH4+) or nitrate (NO3-), although they can also absorb small amounts of amino acids and urea.
Some plants have established symbiotic relationships with fungi and procations capable of reducing atmospheric nitrogen to ammonium, in exchange for which they receive energy molecules from the host plant. The reduced nitrogen is thus incorporated into the food chain (see also the nitrogen cycle). Perhaps the best known case is that of bacteria of the genus Rhizobium with legumes, but there are also associations with bacteria of the genus Frankia and even some cyanobacteria. Finally, some fungi, called ectomycorrhizal fungi, also extend their filaments beyond the reach of the roots, forming mycorrhizae that make more efficient the absorption of nitrites, nitrates and ammonium from the soil in limiting environments.
Isotopes
There are two stable isotopes of nitrogen, 14N and 15N, with the former —produced in the carbon-nitrogen cycle of stars— being the most common without a doubt (99.634%). Ten isotopes of nitrogen have been synthesized, of which 13N has a half-life of nine minutes. In the rest of the isotopes, the half-life is less than 10 seconds, and some isotope even has a half-life much less than one second.
The biological reactions of nitrification and denitrification have a decisive influence on the dynamics of nitrogen in the soil, almost always producing an enrichment of N-15 in the substrate.
Precautions
Nitrogenous fertilizers are an important source of soil and water pollution. Compounds containing cyanide ions form extremely toxic salts and are deadly to many animals, including mammals. Ammonia is also highly toxic.
Effects of Nitrogen on Health
Nitrogen molecules, in their natural state, are found mainly in the air. Nitrogen can be found in water and soil in compounds, in the form of nitrates and nitrites.
Humans have radically changed the natural ratios of nitrates and nitrites, largely due to the application of nitrate-containing manures. Nitrogen is emitted in large quantities by industries. Throughout history, an increase in the presence of nitrates and nitrites in the soil and water has been noted as a consequence of reactions that take place in the nitrogen cycle. This is reflected in an increase in the concentration of nitrogen in the sources used for human consumption, and therefore also in drinking water.
Nitrates and nitrites are known to cause various effects on human health. These are the most common effects:
- It has reactions with hemoglobin in the blood, causing a decrease in oxygen transport capacity by the blood. (nitrito)
- It causes a decrease in the functioning of the thyroid gland. (nitrate)
- It causes low storage of vitamin A. (nitrate)
- It favors the production of nitrosamines, which are known as one of the most common causes of cancer. (nitrates and nitrites)
From a metabolic point of view, nitrogen oxide (NO) is much more important than nitrogen. In 1987, Salvador Moncada discovered that this was a vital messenger in the body for muscle relaxation, and today it is known to be involved in the cardiovascular system, the immune system, the central nervous system, and the peripheral nervous system. The enzyme that produces nitric oxide, nitric oxide synthase, is abundant in the brain.
Although nitric oxide is relatively short-lived, it can diffuse across membranes to carry out its functions. In 1991, a team led by K. E. Anderson at Lund University Hospital, Sweden, demonstrated that nitric oxide triggers an erection by relaxing the muscle that controls blood flow to the penis. The drug Viagra works by releasing nitric oxide to produce the same effect.
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