The non-metallic element of the 15th group [Va] of the periodic table — nitrogen, two atoms of which, when combined, form a molecule — is a colorless, odorless and tasteless gas that makes up most of the Earth’s atmosphere and is an integral part of all life.
Discovery history
Nitrogen gas makes up about 4/5 of the Earth’s atmosphere. It was isolated during early air research. In 1772, the Swedish chemist Karl-Wilhelm Scheele was the first to demonstrate what nitrogen is. In his opinion, air is a mixture of two gases, one of which he called "fiery air", because it supported combustion, and the other - "unclean air", because it remained after the first was consumed. It was oxygen and nitrogen. Around the same time, nitrogen was released by the Scottish botanist Daniel Rutherford, who first published his findings, as well as the British chemist Henry Cavendish and the British clergyman and scholar Joseph Priestley, who shared the primacy of oxygen discovery with Scheele. Further studies showed that the new gas is part of nitrate, or potassium nitrate (KNO 3 ), and, accordingly, it was called nitrogen ("giving saltpeter") by the French chemist Chaptal in 1790. Nitrogen was first assigned to the chemical elements of Lavoisier, whose explanation of the role of oxygen in combustion disproved the theory of phlogiston - popular in the 18th century. erroneous notion of burning. The inability of this chemical element to sustain life (in Greek ζωή) caused Lavoisier to call the gas nitrogen.
The emergence and spread
What is nitrogen? According to the prevalence of chemical elements, it takes sixth place. The atmosphere of the Earth is 75.51% by weight and 78.09% by volume consists of this element and is its main source for industry. The atmosphere also contains a small amount of ammonia and ammonium salts, as well as nitrogen oxides and nitric acid formed during thunderstorms, as well as in internal combustion engines. Free nitrogen is found in many meteorites, volcanic and mine gases and some mineral springs, in the sun, in stars and nebulae.
Nitrogen is also found in mineral deposits of potassium and sodium nitrate, but they are not enough to meet human needs. Another material rich in this element is guano, which can be found in caves where there are many bats, or in dry places visited by birds. Nitrogen is also contained in rain and soil in the form of ammonia and ammonium salts, and in sea water in the form of ammonium ions (NH 4 + ), nitrites (NO 2 - ) and nitrates (NO 3 - ). On average, it makes up about 16% of complex organic compounds, such as proteins, present in all living organisms. Its natural content in the earth's crust is 0.3 parts per 1000. The prevalence in space is from 3 to 7 atoms per silicon atom.
The largest nitrogen producing countries (in the form of ammonia) at the beginning of the 21st century were India, Russia, the USA, Trinidad and Tobago, Ukraine.
Commercial production and use
Industrial nitrogen production is based on fractional distillation of liquefied air. Its boiling point is -195.8 ° C, which is 13 ° C lower than that of oxygen, which is thus separated. 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 residual nitrogen. On a small scale, pure nitrogen is produced by heating barium azide Ba (N 3 ) 2 . Laboratory reactions include heating a solution of ammonium nitrite (NH 4 NO 2 ), the oxidation of ammonia with an aqueous solution of bromine or heated copper oxide :
- NH 4 + + NO 2 - → N 2 + 2H 2 O.
- 8NH 3 + 3Br 2 → N 2 + 6NH 4 + 6Br - .
- 2NH 3 + 3CuO → N 2 + 3H 2 O + 3Cu.
Elemental nitrogen can be used as an inert atmosphere for reactions requiring the exclusion of oxygen and moisture. Finds application and liquid nitrogen. Hydrogen, methane, carbon monoxide, fluorine and oxygen are the only substances that, at the boiling point of nitrogen, do not transform into a solid crystalline state.
In the chemical industry, this chemical element is used to prevent oxidation or other spoilage of the product, as an inert diluent of reactive gas, to remove heat or chemicals, and also as a fire or explosion inhibitor. In the food industry, nitrogen gas is used to prevent spoilage of products, and liquid gas is used for freeze drying and in cooling systems. In the electrical industry, gas prevents oxidation and other chemical reactions, creates pressure in the cable sheath and protects electric motors. In metallurgy, nitrogen is used in welding and brazing, preventing oxidation, carbonization and decarburization. As an inactive gas, it is used in the production of porous rubber, plastic, and elastomers, it serves as a propellant in aerosol cans, and also creates pressure of liquid fuel in jet aircraft. In medicine, rapid freezing with liquid nitrogen is used to preserve blood, bone marrow, tissues, bacteria and sperm. He found application in cryogenic studies.
Connections
Most nitrogen is used in the production of chemical compounds. The triple bond between the atoms of the element is so strong (226 kcal per mole, double that of molecular hydrogen) that the nitrogen molecule hardly enters other compounds.
The main industrial method for fixing an element is the Haber-Bosch ammonia synthesis process developed during World War I to reduce Germany’s dependence on Chilean nitrate. It involves the direct synthesis of NH 3 , a colorless gas with a pungent, irritating odor, directly from its elements.
Most of the ammonia is converted to nitric acid (HNO 3 ) and nitrates - salts and esters of nitric acid, soda ash (Na 2 CO 3 ), hydrazine (N 2 H 4 ) - a colorless liquid used as rocket fuel and in many industrial processes.
Nitric acid is another major commercial compound of this chemical element. A colorless, highly corrosive liquid is used in the manufacture of fertilizers, dyes, medicines and explosives. Ammonium nitrate (NH 4 NO 3 ) - a salt of ammonia and nitric acid - is the most common component of nitrogen fertilizers.
Nitrogen + Oxygen
Nitrogen forms a number of oxides with oxygen, including nitrous oxide (N 2 O), in which its valency is +1, oxide (NO) (+2) and dioxide (NO 2 ) (+4). Many nitrogen oxides are extremely volatile; they are the main sources of pollution in the atmosphere. Nitrous oxide, also known as laughing gas, is sometimes used as an anesthetic. When inhaled, it causes mild hysteria. Nitric oxide quickly reacts with oxygen to form brown dioxide, an intermediate in the production of nitric acid and a powerful oxidizing agent in chemical processes and rocket fuel.
Some nitrides formed by combining metals with nitrogen at elevated temperatures are also used. Boron, titanium, zirconium and tantalum nitrides have special applications. One crystalline form of boron nitride (BN), for example, is not inferior to diamond in hardness and poorly oxidized, therefore it is used as a high-temperature abrasive.
Inorganic cyanides contain the group CN - . Hydrogen cyanide, or HCN hydrocyanic acid , is an extremely unstable and extremely toxic gas that is used for fumigation, ore concentration, and other industrial processes. Dizian (CN) 2 is used as an intermediate chemical and for fumigation.
Azides are compounds that contain a group of three nitrogen atoms —N 3 . Most of them are unstable and very sensitive to shock. Some of them, such as lead azide Pb (N 3 ) 2 , are used in detonators and capsules. Azides, like halogens, readily interact with other substances and form many compounds.
Nitrogen is part of several thousand organic compounds. Most of them are derived from ammonia, hydrogen cyanide, cyanide, nitrous oxide or nitric acid. Amines, amino acids, amides, for example, are derived from ammonia or are closely related to it. Nitroglycerin and nitrocellulose are esters of nitric acid. Nitrites are obtained from nitrous acid (HNO 2 ). Purines and alkaloids are heterocyclic compounds in which nitrogen replaces one or more carbon atoms.
Properties and reactions
What is nitrogen? It is a colorless, odorless gas that condenses at -195.8 ° C to a colorless, low-viscosity liquid. The element exists in the form of N 2 molecules, represented in the form: N ::: N:, in which the binding energy of 226 kcal per mole is second only to carbon monoxide (256 kcal per mole). For this reason, the activation energy of molecular nitrogen is very high, therefore, under normal conditions, the element is relatively inert. In addition, the highly stable nitrogen molecule contributes significantly to the thermodynamic instability of many nitrogen-containing compounds in which the bonds, although strong enough, are inferior to molecular nitrogen bonds.
Relatively recently and unexpectedly, the ability of nitrogen molecules to serve as ligands in complex compounds was discovered. The observation that some solutions of ruthenium complexes can absorb atmospheric nitrogen has led to the fact that soon a simpler and better way to fix this element can be found.
Active nitrogen can be obtained by passing low-pressure gas through a high voltage electrical discharge. The product glows with yellow light and is much more willing to react than molecular with atomic hydrogen, sulfur, phosphorus and various metals, and is also able to decompose NO to N 2 and O 2 .
A clearer idea of what nitrogen is can be obtained thanks to its electronic structure, which has the form 1s 2 2s 2 2p 3 . Five electrons of the outer shells weakly shield the charge, as a result of which the effective nuclear charge is felt at a distance of the covalent radius. Nitrogen atoms are relatively small and have a high electronegativity, located between carbon and oxygen. The electronic configuration includes three semi-filled outer orbitals, allowing the formation of three covalent bonds. Therefore, the nitrogen atom must have extremely high reactivity, forming stable binary compounds with most other elements, especially when the other element is significantly different in electronegativity, which gives a significant polarity to the bonds. When the electronegativity of the other element is lower, the polarity gives the nitrogen atom a partial negative charge, which releases its undivided electrons to participate in coordination bonds. When the other element is more electronegative, a partially positive nitrogen charge substantially limits the donor properties of the molecule. With a small polarity of the bond, due to the equal electronegativity of the other element, multiple bonds prevail over single bonds. If the mismatch in atomic dimensions prevents the formation of multiple bonds, then the formed simple bond is likely to be relatively weak, and the connection will be unstable.
Analytical chemistry
Often the percentage of nitrogen in a gas mixture can be determined by measuring its volume after absorption of other components by chemical reagents. The decomposition of nitrates with sulfuric acid in the presence of mercury releases nitric oxide, which can be measured as a gas. Nitrogen is released from organic compounds when they burn over copper oxide, and free nitrogen can be measured as gas after absorption of other combustion products. The well-known Kjeldahl method for determining the content of the substance under consideration in organic compounds consists in the decomposition of the compound with concentrated sulfuric acid (if necessary, containing mercury or its oxide, as well as various salts). Thus, nitrogen is converted to ammonium sulfate. Adding sodium hydroxide releases ammonia, which is collected with ordinary acid; the residual amount of unreacted acid is then determined by titration.
Biological and physiological significance
The role of nitrogen in living matter confirms the physiological activity of its organic compounds. Most living organisms cannot use this chemical element directly and must have access to its compounds. Therefore, nitrogen fixation is of great importance. In nature, this occurs as a result of two main processes. One of them is the effect of electrical energy on the atmosphere, due to which the nitrogen and oxygen molecules dissociate, which allows free atoms to form NO and NO 2 . The dioxide then reacts with water: 3NO 2 + H 2 O → 2HNO 3 + NO.
HNO 3 dissolves and comes to Earth with rain in the form of a weak solution. Over time, the acid becomes part of the combined soil nitrogen, where it is neutralized, forming nitrites and nitrates. N content in cultivated soils, as a rule, is restored due to the application of fertilizers containing nitrates and ammonium salts. Excretion of animals and plants and their decomposition returns nitrogen compounds to soil and air.
Another main process of natural fixation is the activity of legumes. Due to symbiosis with bacteria, these cultures are able to convert atmospheric nitrogen directly into its compounds. Some microorganisms, such as Azotobacter Chroococcum and Clostridium pasteurianum, are capable of fixing N independently.
The gas itself, being inert, is harmless, except when breathing it under pressure, and it dissolves in blood and other body fluids in higher concentrations. This causes a narcotic effect, and if the pressure decreases too quickly, an excess of nitrogen is released in the form of gas bubbles in various places of the body. It can cause muscle and joint pain, fainting, partial paralysis, and even death. These symptoms are called decompression sickness. Therefore, those who are forced to breathe air under such conditions must very slowly reduce the pressure to normal so that excess nitrogen exits through the lungs without the formation of bubbles. A better alternative is to use a mixture of oxygen and helium for breathing. Helium is much less soluble in body fluids, and the danger is reduced.
Isotopes
Nitrogen exists in the form of two stable isotopes: 14 N (99.63%) and 15 N (0.37%). They can be separated by chemical exchange or by thermal diffusion. The nitrogen mass in the form of artificial radioactive isotopes is in the range of 10-13 and 16-24. The most stable half-life of 10 minutes. The first artificially induced nuclear transmutation was performed in 1919 by the British physicist Ernest Rutherford, who, by bombarding nitrogen-14 with alpha particles, received oxygen-17 nuclei and protons.
The properties
Finally, we list the main properties of nitrogen:
- Atomic number: 7.
- Atomic mass of nitrogen: 14.0067.
- Melting point: -209.86 ° C.
- Boiling point: -195.8 ° C.
- Density (1 atm, 0 ° C): 1.2506 g of nitrogen per liter.
- Typical oxidation states: -3, +3, +5.
- Electron Configuration: 1s 2 2s 2 2p 3 .