Nitrogen (N), a non-metallic element of group 15 [Va] from the periodic table. It is a colorless, odorless, and tasteless gas that is the most abundant element in the Earth’s atmosphere and is the constituent substance of all living things.
About four-fifths of the Earth’s atmosphere is nitrogen, Which in early research isolated air and was identified as a special substance. Carl Wilhelm Schell, a Swedish chemist, showed in 1772 that air is a mixture of two gases, one of which he calls “fire air” because it supports combustion, and the other “unclean air”, because after It remained. “Fire air” was used, but “fire air” was oxygen and nitrogen was “unclean air”. Around the same time, nitrogen was introduced by a Scottish botanist, Daniel Rutherford (who was the first to publish his findings), by the English chemist Henry Cavendish, and the English clergyman and scientist Joseph Priestley, who, along with Shell, Was recognized. Credits are given for the discovery of oxygen. Subsequent work showed that the new gas as the constituent of nitrate was the common name for potassium nitrate (KNO3), and accordingly, it was called the French chemotherapy of Jean-Antoine en Claude Sheptal in 1790 from nitrogen. Nitrogen was first considered a chemical. An element by Antoine Laurent Lavoisier, who explained the role of oxygen in combustion, ultimately overturned Flogouston’s theory, a misconception of combustion that became common in the early 18th century. The inability of nitrogen to support life (Greek: zoe) led Lavoisier to call it nitrogen, which is still the French equivalent of nitrogen.
Event and distribution
Among the elements, nitrogen ranks sixth in cosmic abundance. The Earth’s atmosphere is composed of 75.51% by weight (or 78.09% by volume) of nitrogen. It is the main source of nitrogen for trade and industry. Barley contains slightly different amounts of ammonia and ammonium salts, as well as nitrogen oxides and nitric acid (the latter formed in electrical storms and in internal combustion engines). Free nitrogen is found in many meteorites. In volcanic gases, mines and some mineral springs. In the sun; And in some stars and nebulae
Nitrogen also occurs in the mineral reserves of nitrate or salt (potassium nitrate, KNO3) and Chilean salt (sodium nitrate, NaNO3), but these reserves are present in amounts that are insufficient for human needs. Another nitrogen-rich substance is gano, which is found in bat caves and in dry places where birds are used. In combination, nitrogen is found in rain and soil as ammonia and ammonium salts and in seawater as ammonium (NH4 +), nitrate (NO2-) and nitrate (NO3-) ions. Nitrogen, on average, makes up about 16 percent of the weight of complex organic compounds known as proteins that are present in all living organisms. The natural abundance of nitrogen in the Earth’s crust is 0.3 parts per 1000. Cosmic abundance – an estimate of the total abundance in the universe – is between three and seven atoms per silicon atom, which is considered 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.
Production and commercial use
Commercial production of nitrogen is mainly done by distillation of liquid air fraction. The boiling point of nitrogen is 95195.8 ° C (-320.4 ° F), about 13 ° C (-23 ° F) lower than oxygen, which is why it is discarded. Nitrogen can also be produced on a large scale by burning carbon or hydrocarbons in air and separating carbon dioxide and water from the remaining nitrogen. On a small scale, pure nitrogen is made by heating barium azide, Ba (N3) 2. Various laboratory reactions involving nitrogen from nitrogen include heating solutions of ammonium nitrate (NH4NO2), oxidation of ammonia by bromine, and oxidation of ammonia by gram-oxide.
Nitrogen can be used as an inert space for reactions that require the release of oxygen and moisture. In liquid form, nitrogen has valuable applications. With the exception of hydrogen, methane, carbon monoxide, fluorine, and oxygen gases, almost all chemicals have a low vapor pressure at the nitrogen boiling point and therefore exist as crystalline solids at that temperature.
In the chemical industry, nitrogen is used as a substance to prevent oxidation or spoilage of another product, as an inert diluent of a reactive gas, as a carrier for the removal of heat or chemicals, and as a fire or explosion inhibitor. In the food industry, nitrogen gas is used to prevent wastage through oxidation, mold or insects, and liquid nitrogen is used to dry ice and refrigeration systems. In the electrical industry, nitrogen is used to prevent oxidation and other chemical reactions, to pressurize cable coats, and to protect motors. Nitrogen is used in the metal industry in welding, soldering and turning, where it helps prevent oxidation, carbide and desorption. As a non-reactive gas, nitrogen is used to make rubber, plastics and elastomers as a fuel propellant, as a combustion gas for aerosol cans and to pressurize liquid fuels for reaction jets. In medicine, rapid freezing with liquid nitrogen may be used to preserve blood, bone marrow, tissue, bacteria, and semen. Liquid nitrogen is also useful in cryogenic research.
Although other uses are important, most of the nitrogen is used to make nitrogen compounds. The triple bond between atoms in nitrogen molecules is so strong (226 kcal per mole, more than twice as much as molecular hydrogen) that it is difficult to enter.
The main commercial method for nitrogen modification (nitrogen composition of elements in compounds) is the Haber-Bush process for ammonia synthesis. This trend was developed during World War I to reduce Germany’s dependence on Chilean nitrate. This involves the direct synthesis of ammonia from its elements.
Large amounts of nitrogen along with hydrogen are used to produce ammonia, NH3, a colorless gas with a pungent, irritating odor. The main trade method for ammonia synthesis is the Haber-Bush process. Ammonia is one of the two major nitrogen compounds in the trade. This substance has many applications in making other important nitrogen compounds. Much of the commercially synthesized ammonia is converted to nitric acid (HNO3) and nitrates, which are salts and esters of nitric acid. Ammonia in the soda process Ammonia (Solvay process) is used to make soda ash, Na2CO3. Ammonia is also used in the preparation of hydrazine, N2H4, a colorless liquid as rocket fuel and in many industrial processes.
Nitric acid is another popular commercial compound of nitrogen. A colorless, highly corrosive liquid, widely 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 nitrogen component of synthetic fertilizers and is used with fuel oil as an explosive (ANFO).
With oxygen, nitrogen forms several oxides, including nitrogen oxide, N2O, in which nitrogen is in the +1 oxidation state. Nitric oxide, NO, where it is in the +2 state. And nitrogen dioxide, NO2, where it is in the +4 state. Many nitrogen oxides are very unstable. They are the main sources of atmospheric pollution. Nitric oxide, also known as laughing gas, is sometimes used as an anesthetic. Causes mild hysteria when inhaled. Nitric oxide reacts rapidly with oxygen to form brown nitrogen dioxide, a mediator in the production of nitric acid and a powerful oxidant used in chemical processes and rocket fuels.
For some nitrides, solids that are formed by the direct combination of metals with nitrogen, usually at high temperatures, are of particular importance. These materials include hardeners produced when alloy steels are heated in the atmosphere, a process called nitration. Those with boron, titanium, zirconium and tantalum have special applications. For example, a crystalline form of boron nitride (BN) is about as hard as diamond and is easily oxidized and is useful as an abrasive at high temperatures.
Mineral cyanides contain the CN− group. Hydrogen cyanide or fermionetrile, HCN, is a highly volatile and highly toxic gas used in steam, rock concentrations, and other industrial processes. Cyanogen or oxalonitrile (CN) 2 is also used as an intermediate and heating chemical.
Azides, which may be inorganic or organic, are compounds consisting of three nitrogen atoms as a group, denoted as (3N3). Most azides are unstable to shock and are very sensitive. Some of them, such as lead azide, lead (N3) 2 are used in explosives and percussion caps. Azides, like halogen compounds, react easily with other substances by moving the so-called azide group, producing different types of compounds.
Nitrogen forms thousands of organic compounds. Many known types may be considered ammonia, hydrogen cyanide, cyanogen, and nitrogen or nitric acid. For example, amines, amino acids and amides are derived from ammonia. Nitroglycerin and nitrocellulose are nitric acid esters. Nitro compounds are formed by the reaction (called nitration) between nitric acid and an organic compound. Nitrites are derived from nitric acid (HNO2). Nitrosu compounds are obtained by the action of nitric acid on an organic compound. Purines and alkaloids are heterocyclic compounds in which nitrogen replaces one or more carbon atoms.
Properties and reactions
Nitrogen is a colorless, odorless gas that condenses into a colorless, mobile liquid at 951 ° C. This element exists as N2 molecules, such as: N ::: N:, for which the bond energy of 226 kcal / mol with carbon monoxide alone exceeds 256 kcal / mol. Due to this high bond energy, the activation energy for the molecular nitrogen reaction is usually very high, making nitrogen relatively inert in most reagents under normal conditions. In addition, the high stability of the nitrogen molecule contributes to the thermodynamic instability of many nitrogen compounds, in which the bonds, although reasonably strong, are far lower than those in the molecular nitrogen. For these reasons, nitrogen seems to obscure the quite effective nature of the reactivity of its individual atoms.
A relatively new and unexpected discovery is that nitrogen molecules are able to act as ligands in complex compounds. The observation that some solutions of the ruthenium complex can absorb atmospheric nitrogen has led me to hope that one day a simpler and better way of modifying nitrogen will be found.
An active form of nitrogen, possibly containing free nitrogen atoms, can be created by passing low-pressure nitrogen gas through a high-pressure electrical discharge. This product is bright with yellow light and is much more reactive than ordinary molecular nitrogen and combines with atomic hydrogen and sulfur, phosphorus and various metals and has the ability to decompose nitric oxide, NO, N2 and O2.
A nitrogen atom has an electronic structure represented by 1s22s22p3. The five outer shell electrons show a very weak nuclear charge, with the result that the effective nuclear charge felt at a distance of the covalent radius is relatively high. Therefore, nitrogen atoms are relatively small in size and high in energy, and mediate between carbon and oxygen in both properties. The electronic configuration consists of three semi-solid outer orbits that enable the atom to form three covalent bonds. Therefore, the nitrogen atom must be a highly reactive species and combine with most other elements to form stable binary compounds, especially when the other element is different enough in electronegativity to transfer significant polarity to the bonds. When the other element is less powerful than nitrogen, the polarity gives a slight negative charge to the nitrogen atom, leaving the pair electrons only available for coordination. However, when the other element is more electro-nutritious, however, the partial positive charge due to nitrogen greatly limits the donor property of the molecule. When the bond polarity is low (due to the electron being of another element that is similar to nitrogen), multiple bonds are strongly favored over single bonds. If the atomic size difference prevents such multiple bonds, the single bond that forms is likely to be relatively weak, and the combination is likely to be unstable due to the free elements. All these properties of nitrogen bond are visible in its general chemistry However, when the other element is more electro-nutritious, however, the partial positive charge due to nitrogen greatly limits the donor property of the molecule. When the bond polarity is low (due to the electron being of another element that is similar to nitrogen), multiple bonds are strongly favored over single bonds. If the atomic size difference prevents such multiple bonds, the single bond that forms is likely to be relatively weak, and the combination is likely to be unstable due to the free elements. All these properties of nitrogen bond are visible in its general chemistry However, when the other element is more electro-nutritious, however, the partial positive charge due to nitrogen greatly limits the donor property of the molecule. When the bond polarity is low (due to the electron being of another element that is similar to nitrogen), multiple bonds are strongly favored over single bonds. If the atomic size difference prevents such multiple bonds, the single bond that forms is likely to be relatively weak, and the combination is likely to be unstable due to the free elements. All these properties of nitrogen bond are visible in its general chemistry If the atomic size difference prevents such multiple bonds, the single bond that forms is likely to be relatively weak, and the combination is likely to be unstable due to the free elements. All these properties of nitrogen bond are visible in its general chemistry If the atomic size difference prevents such multiple bonds, the single bond that forms is likely to be relatively weak, and the combination is likely to be unstable due to the free elements. All these properties of nitrogen bond are visible in its general chemistry
Often the percentage of nitrogen in gaseous mixtures can be determined by measuring the volume after the absorption of other components by chemicals. Decomposition of nitrate by sulfuric acid in the presence of mercury releases nitric oxide, which can be measured as a gas. Nitrogen is released from organic compounds when burned on copper oxide, and after absorbing other combustion products, free nitrogen can be measured as a gas. Kjeldahl’s famous method for determining the nitrogen content of organic compounds involves the digestion of compounds with concentrated sulfuric acid (optional containing mercury, or its oxides and various salts, depending on the nature of the nitrogen composition). In this way, the available nitrogen is converted to ammonium sulfate. In addition it releases an excess of free sodium ammonia hydroxide, which accumulates in standard acid. The amount of residual acid, which did not react with ammonia, is determined by titration.
Biological and physiological importance
As may be due to the importance of the presence of nitrogen in living matter, most organic nitrogen compounds are physiologically active. Most organisms cannot use nitrogen directly and must have access to its compounds. Therefore, nitrogen fixation is vital. In nature, two main processes of nitrogen fixation are known. One is the action of electrical energy on the atmosphere, which separates nitrogen and oxygen molecules and allows free atoms to form nitric oxide, NO, and nitrogen dioxide, NO2. Nitrogen dioxide then reacts with water.
It is as follows :
Nitric acid, HNO3, dissolves as a very dilute solution and comes to the ground with rain. Over time, it is part of the soil nitrogen compound, where it is neutralized, converted to nitrate and nitrate. The amount of nitrogen in the soil under cultivation is artificially enriched and renewed by fertilizers containing nitrate and ammonium salt. Animal and plant decomposition returns nitrogen compounds to the soil and air, and some bacteria in the soil break down nitrogen compounds and return the element to the air.
The main process of modifying the natural nitrogen of some plants and vegetables is called legumes. Through a collaboration with bacteria, legumes are able to convert atmospheric nitrogen directly into nitrogen compounds. Some bacteria alone, such as Azotobacter chroococcum and Clostridium pasteurianum, are also able to remove nitrogen.
Nitrogen is ineffective in itself, except when it is breathed under pressure, it is innocent and in this case it dissolves more than normal concentration in blood and other body fluids. This in itself has a narcotic effect, but if the pressure drops too quickly, the excess nitrogen will evaporate as a gas bubble in different parts of the body. These can cause muscle and joint pain, fainting, minor weakness and even death. These symptoms are called “bending” or decompression sickness. Divers, aeronauts, those working in deep caissons where the air pressure drops too fast, and others who have to breathe air are under pressure so they must be very careful not to let the pressure drop too slowly following exposure. . This allows excess nitrogen to be released safely through the lungs without creating bubbles. A better alternative is to replace oxygen and helium with air.
Nitrogen exists as two stable isotopes, 14N (frequency 99.63%) and 15N (frequency 0.37%). These can be separated by chemical exchange or heat dissipation. Synthetic radioactive isotopes have masses of 10-13 and 24-24. The most stable of them has a half-life of only about 10 minutes. The first artificially induced nuclear transfer was reported by the English physicist Ernest Rutherford, who bombarded nitrogen 14 with alpha particles to bombard 17 oxygen nuclei and protons (1919).
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