Nitrogen

Nitrogen, chemical element, symb. N (Fr. AZ), atomic number 7, at. wt. 14.008; boiling point -195.7°; 1 liter of N at 0° and 760 mm pressure weighs 1.2508 g [Lat. Nitrogenium («nitre-producing»), Germ. Stickstoff («suffocating substance»), Fr. Azote (from Gr. a-negation, zoe-life), i.e., «lifeless»]; belongs to Group V of the periodic system. N can form five-valence electropositive and three-valence electronegative ions. Under the influence of bombardment by α-particles, flying away during the radioactive decay of the radioactive element Ra-C, atoms of N were «shot» in 1919 by Rutherford and Chadwick, with the formation of atoms of element H as fragments of destroyed N atomic nuclei. Chemical properties. N is a colorless gas, slightly soluble in water: at 0° only 2.35 vol. %, at 10° -1.8%, at 38° -1.22%. In blood solubility is even less - at 38° only 1.1 vol. %. However, at low partial pressure of N, its content in blood is still relatively somewhat higher, due to the fact that N is unstably bound with blood colloids, mainly with blood pigment. N under ordinary conditions of temperature and pressure is an extremely inert element, not entering into chemical reaction with other elements. This explains the presence in nature of significant amounts of free, not chemically combined with other elements, N, of which ⁴/₅ by volume consists the atmosphere of the Earth. (The total content of nitrogen in the entire earth's crust does not exceed 0.02%, including the atmosphere.) Although living organisms cannot exist in an atmosphere of N, on Earth there is no living creature, plant or animal, in the organism of which N would not play an essential role. But this role is played only by N already in the state of chemical combination with other elements: free N neither plants (with the exception of a few species of microorganisms) nor animals can assimilate. Besides its presence in the atmosphere, nitrogen occurs in nature, mainly in the form of nitrogen compounds, in particular proteins, forming the most essential component of protoplasm. In the process of metabolism in the animal organism, protein compounds undergo a series of destructive transformations, as a result of which the N contained in them is excreted in the form of urea, etc. Urea, as well as other products of decomposition of nitrogenous substances, formed during the decay of animal and plant residues, under the influence of the life activity of microorganisms in the soil are converted into ammonia, which is further oxidized into nitrous and nitric acids, forming in the soil corresponding salts (nitrate and nitrite salts). The result of such micro-biological processes are natural deposits of nitrate salts (nitrates). The same processes form the basis for the artificial production of nitrogen salts in «niter beds» (salpetriere), in which wastes containing N are subjected to biological processes, ultimately yielding nitre. In a dry climate, preventing the leaching of nitrate salts from the soil, they can accumulate in it in huge quantities. This is the origin of the famous natural deposits of sodium nitrate in Chile, for a long time one of the main sources of combined N necessary for fertilizers, for the production of explosives, for aniline dyes, medicines, and many other products of the chemical industry. Fixation of atmospheric nitrogen. The end of the 19th century and the past part of the 20th century have been marked by the greatest achievements of chemical technology and the victory of man over nature in the field of fixation of atmospheric N on a technical scale. The latter is carried out mainly by four methods: 1) the Birkeland and Eyde method, repeating the process occurring in nature during atmospheric electrical discharge; it consists in the fact that air is blown through the flame of an alternating or direct current electric arc (Schönberger), stretched under the action of a strong electromagnet into a wide disk 2 m in diameter. As a result, the products of the endothermic reaction (N₂ + O₂=2NO), which occurs quite quickly at the temperature of the electric arc, are quickly removed from it and cooled to a temperature of about 1500°. At this temperature the reverse reaction (2NO = N₂ + O₂) already has a significantly lower rate. Further, 2NO is easily oxidized by air to 2NO₂=N₂O₄. Coming into contact with water, N₂O₄ forms a mixture of HNO₂ and HNO₃. Under certain conditions, it is possible to obtain mainly HNO₃, which is either concentrated by distillation with H₂SO₄ or neutralized with lime to form Ca(NO₃)₂, calcium or «Norwegian» saltpeter. The process of formation of HNO₃ is economically associated with the expenditure of a large amount of cheap hydroelectric energy necessary to create a powerful electric arc. Therefore, it is profitable only in the presence of rich hydroelectric resources in a given area, waterfalls, etc. 2) The cyanamide method of Frank and Caro is based on the fact that calcium carbide CaC₂ at high temperature is capable of attaching N₂ to itself with the formation of calcium cyanamide CaCN₂ according to the reaction: CaC₂+N₂=CaCN₂+C. Calcium cyanamide is either directly used as fertilizer, decomposing in the soil with water to form ammonia, or is treated in autoclaves with water vapor, forming NH₃, neutralized in special «saturators» with sulfuric or nitric acids. 3) The method of obtaining synthetic ammonia (Haber and Bosch) from elements is based on the fact that at temperatures of 500-600° and a pressure of 200 atm, nitrogen and hydrogen under the influence of catalysts (osmium and other rare metals) combine exothermically to form ammonia with the release of such an amount of heat that it itself maintains the course of the already started reaction. The Haber-Bosch method largely resolved the problem of the «nitrogen famine» of humanity, which was looming in connection with the approaching final depletion of Chilean nitrate deposits and the increase in the population of the Earth, and with it the increase in the world's need for nitrogen products for the chemical industry. 4) The Serp method: by heating aluminum oxide and coke at 1500-1700° in N, AlN is obtained; the latter with water gives ammonia: AlN+3H₂O = NH₃+Al(OH)₃. Application of N. As an inert gas, N is used in medicine for immobilization of affected areas of the lungs in tuberculosis (pneumothorax), in technology - for protecting metals from the action of active gases, for inflating automobile tires, which change at high pressure from the action of air, for preventing fires during the transfer of fuel, etc., but mainly for the production of calcium cyanamide and synthetic ammonia. Almost pure N (with only about 3% argon impurity) is obtained by fractional distillation of liquid air; for the Haber cyanamide synthesis, N in mixture with H₂ can be obtained from flue gases after absorption of CO₂, H₂O and residual O₂ from them, since most of the atmospheric oxygen in them has already combined with carbon and hydrogen. In the laboratory, N in mixture with argon (see) is obtained from air by absorption of O, water vapor and CO₂ from it. Completely pure N is obtained by decomposition of ammonium nitrite salt: NH₄NO₂→2H₂O+N₂. Active N becomes either at high temperature or after passing strong electric sparks through it. Such «active» N reacts with Na, Zn, Hg, Pb, Sn, P, forming nitrides, but, for still unexplained reasons, does not react with either O or H. Compounds of nitrogen. The simplest compounds of N are its compounds with O, H and C. Of the first, five oxides of N should be named: 1) nitrous oxide, N₂O, a colorless gas supporting combustion, also called «laughing gas». Obtained by heating ammonium nitrate according to the equation: NH₄NO₃=2H₂O+N₂O or by reduction of higher oxides of N. Used in medicine as an anesthetic agent for tooth extraction and other brief surgical operations. 2) Nitric oxide, NO, a colorless gas, quickly browning in air due to the addition of O and conversion to nitrogen dioxide NO₂, which in turn partially polymerizes to N₂O₄ according to the equation: 2NO₂=N₂O₄. At high temperatures the equilibrium shifts to the left with the formation in predominant amount of NO₂ molecules. 3) Nitrous anhydride, N₂O₃, is formed by heating nitric acid in the presence of starch and is at low temperature a blue liquid, easily decomposing at room temperature to NO and N₂O₄.

Nitrogen anhydride with water gives nitrous acid: N₂O₃ + H₂O = 2HNO₂, which is also formed in the soil under the influence of certain microorganisms; therefore the presence of nitrous acid in drinking water should be considered a sign of its contamination. 4) Tetroxide of N., N₂O₄, is formed by the oxidation of NO by atmospheric oxygen and by the decomposition of lead nitrate: Pb(NO₃)₂ = PbO + O + N₂O₄. It is a liquid that boils at 25-26° with the evolution of brown fumes. When N₂O₄ is dissolved in water, nitric and nitrous acids are formed: N₂O₄ + H₂O = HNO₃ + HNO₂. 5) Nitrogen anhydride, N₂O₅, gives with water nitric acid HNO₃; it is a solid crystalline substance with strongly oxidizing properties. Nitric acid is obtained either from saltpeter or from the air by the Birkeland-Eyde method or by the catalytic oxidation of ammonia with air under the catalytic action of platinum (the Ostwald-Kühlmann method).-All oxides of N. are extremely poisonous substances. Oxides of nitrogen as a professional hazard. The oxides of nitrogen found in factory air are a mixture consisting of nitric oxide (NO), nitrogen dioxide (NO₂), nitrous acid (HNO₂), and nitric acid (HNO₃). They are formed from fuming nitric acid when it acts on reducing substances: metals, organic substances, sulfuric acid and its salts. Due to the widespread use of nitric acid in industry, the number of industries and occupations where there is a possibility of the release of nitrogen oxides into the air, and consequently the danger of poisoning by them, is very great. Among the most important industries should be mentioned: the production of nitric acid (from saltpeter and by oxidation of N. from the air), the numerous nitration processes in chemical plants (in the production of coal-tar dyes), the production of explosives (smokeless powder, nitroglycerin, picric acid, tetryl, fulminating mercury and many others,-all these are nitrated substances), the production of celluloid, the production of artificial fertilizers (superphosphates), many metalworking industries, especially the etching of brass and copper products, sulfuric acid plants, mining (blasting operations) and many others. Nitrogen oxides are formed particularly easily and in large quantities during metal etching and nitration processes.-According to Lehmann and Hasegawa, nitrogen oxides should be practically considered as an equimolecular mixture of nitrous and nitric acids, since nitric oxide, however it is formed, is immediately oxidized by atmospheric oxygen and converted into nitrogen dioxide (brown fumes), and the latter with water (in humid air or on human mucous membranes) splits according to the equation: 2NO₂ + 4H₂O = HNO₂ + HNO₃. Poisoning by nitrogen oxides is exclusively acute in nature. Chronic poisonings are unknown. The main action of nitrogen oxides on the body consists in cauterizing the epithelium of the respiratory tract and alveoli with subsequent pulmonary edema. When inhaling gases of high concentration, their absorption and effect on the blood (formation of methemoglobin) occurs.-S-ymptoms: in mild poisonings, the painful phenomena are limited to irritation of the upper respiratory tract, scratching in the throat, cough, and a feeling of tightness in the chest. These phenomena disappear quickly. With a stronger effect on the epithelium of the respiratory tract, the phenomena of actual poisoning are preceded by a certain latent period (10 hours or more), during which the victim, experiencing only mild irritation, feels relatively well. Only gradually does the picture of severe damage to the respiratory tract develop: a feeling of compression of the airways appears, dry cough, gradually increasing shortness of breath, sometimes vomiting, severe cyanosis; viscous bloody sputum begins to be secreted. Consciousness is completely preserved all the time. After 30-40 hours, with increasing shortness of breath (pulmonary edema) and weakening of cardiac activity, death occurs. Recovery is possible, but it is slow. Sometimes severe phenomena of sudden shortness of breath and cyanosis developed after 8-10 days; they can also end in death. The origin of such late lesions is explained by the fact that in the respiratory tract and alveoli, at the affected sites, granulations form during the healing process, causing a severe narrowing of the airways. In very severe cases, after inhaling nitrogen oxides of high concentration, they can be absorbed into the bloodstream and cause the formation of methemoglobin in it. Finally, according to some authors, a direct effect of nitrogen oxides on the respiratory centers of the brain is also possible; death then occurs almost instantaneously. There are very sharp individual differences regarding predisposition to poisoning by nitrogen oxides. There are known cases where from a group of people subjected to the same danger of poisoning, some died, some suffered more mild lesions, and some remained completely healthy. Persons with diseased or weakened respiratory organs are most easily poisoned. Regarding toxic doses, there are data from Lehmann, who gives the following figures: 0.1 mg of nitric acid per liter of air has no harmful effect, 0.2 mg strongly irritates, 0.3-0.5 mg can only be tolerated for a short time; 0.5-0.7 mg causes the death of a cat within 2-4 hours, above 0.9 mg kills it within 1-2 hours.-Treatment. The most effective means is the inhalation of oxygen, which should be carried out until the pulmonary symptoms completely disappear. It is advisable to have water vapor inhaled simultaneously. Inhalation of oxygen should be resorted to during the latent period if poisoning by nitrogen oxides has been established. In other cases, bloodletting is also resorted to. - Preventive measures: hermetic sealing of equipment; carrying out nitration processes in closed apparatus or in fume hoods equipped with proper exhaust ventilation; installation of powerful exhaust ventilation in workrooms. In other rapidly proceeding chemical processes, the wearing of gas masks is necessary. Compounds of nitrogen with hydrogen, in contrast to oxides of N., are exothermic. Among them, the most common is ammonia, NH₃, formed by the decomposition of urea under the influence of microorganisms, and therefore also during the putrefaction of protein substances. Ammonia is released (in the production of illuminating gas) from coal or during the coking of coal; in recent years, NH₃ is obtained directly from the gaseous elements N₂ and H₂ by the Haber-Bosch method. In laboratories, ammonia is usually obtained by the action of slaked lime [Ca(OH)₂] on sal ammoniac (NH₄Cl): 2 NH₄Cl + Ca(OH)₂ = CaCl₂ + 2 NH₃ + 2 H₂O.-Besides ammonia, the following are known: the compound of N. with H of composition N₃H-hydrazoic acid, forming extremely explosive salts-azides; hydroxylamine, NH₂OH, and hydrazine, N₂H₄, which have strongly reducing properties and form salt-like compounds with acids. Ammonia NH₃ in aqueous solution has alkaline properties, with acids it forms so-called ammonium salts, for example, NH₃ + HCl = NH₄Cl-ammonium chloride, or sal ammoniac. In the absence of water, ammonia is a weak acid and forms salts, for example, NH₂Na-sodium amide. When the metal is replaced by a hydrocarbon radical, a primary amine is obtained, for example, CH₃NH₂-methylamine, C₂H₅NH₂-ethylamine, etc. Hydrocarbon radicals can replace not only one, but two and all three H atoms, for example, (C₂H₅)₂NH (secondary amine)-diethylamine, (CH₃)₃N (tertiary amine)-trimethylamine. Compounds are also known which are formed as a result of addition, for example, of halogen derivatives to a tertiary amine, for example, (CH₃)₃N + CH₃I = (CH₃)₄NI. These compounds, analogous to ammonium salts, are called quaternary ammonium bases (in the example given-methyl iodide tetramethylammonium). If the amide group NH₂ enters the radical of an organic acid, then such an acid is called an amino acid, for example, aminoacetic acid, or glycine, CH₂NH₂.COOH. By replacing the hydroxyl group in an organic acid with an amide group, the amide of the acid can be obtained, for example, the amide of acetic acid-CH₃CONH₂.

Very important nitrogenous organic compounds in organic chemistry should be considered nitro compounds, easily obtained by the action of strong nitric acid on aromatic hydrocarbons (e.g., nitrobenzene, C6H5NO2) or somewhat more difficultly by the action of dilute nitric acid, under pressure, on fatty hydrocarbons. A special, very important and far from fully studied class of nitrogenous organic compounds are the alkaloids (see.), partly artificially synthesized, partly extracted from parts of plant or animal organisms. Besides oxygen and hydrogen compounds of N., compounds of N. with C are also known: gaseous and very poisonous dicyan, C2N2, and hydrocyanic, or prussic acid and its salts—the cyanides. Cyanide compounds are formed by fusing with caustic alkali or alkali metals nitrogenous organic compounds (waste from slaughterhouses—hooves, hides, horns, etc.). Cyanide compounds form various complex compounds, of which yellow blood salt and red blood salt should be particularly noted. The largest number of nitrogenous compounds belong to organic nitrogenous compounds. Among them, protein substances (see Proteins) have the most complex composition, representing essentially chains of amino acids (see.) in various combinations. The simplest products of decomposition of protein compounds and intermediate compounds in animal and plant organisms are urea, uric acid, asparagine, guanidine, and others. Nitrogen cycle. The largest amount of N. on our planet is in the free state in the atmosphere. Atmospheric N. consists of N2 molecules, in which the nitrogen atoms are very firmly bound to each other. Dissociation of the N. molecule into atoms and their ionization require the expenditure of a considerable amount of energy. This explains the inertness of free N. But once it has entered into chemical compounds, N. gives them greater reactivity. This property of nitrogenous substances determines the role that they play in the evolution of living matter. Plants draw upon compounds of bound N. from the soil as material for constructing the nitrogenous compounds of their organisms (proteins and other substances). Animals assimilate the proteins synthesized in plants, then in the process of metabolism (see.) subject these protein compounds to destructive decomposition reactions, excreting, as decomposition products, urea, uric acid, creatinine, and other compounds into the soil or into the water of seas and rivers; into the soil or water also passes the entire organism with all the nitrogenous substances contained in it upon its death. In the soil, nitrogenous substances of animal and plant origin undergo decomposition under the action of microorganisms and ultimately turn into ammonia and ammonium salts. After this, other microorganisms oxidize ammonia by means of atmospheric O to nitrous and nitric acid (Winogradsky). The nitrites and nitrates formed in the soil are easily absorbed by plant roots and assimilated through a whole (not yet fully elucidated) series of intermediate compounds (formhydroxamic acid—HO·CH:NOH, asparagine, etc.) with the formation of complex molecules of nitrogen-containing compounds—proteins, polypeptides, alkaloids, etc. Thus, in nature a constant cycle of compounds of bound N. can be observed from plants to animals and back, through soil and water, to plants. Into this cycle, however, as side processes, flow certain processes that fix atmospheric N. and thereby increase the total available reserve of bound N. in nature. Such processes are: 1. Atmospheric electricity discharges, causing reactions of combination of atmospheric N. with O to form oxides of N., which, dissolving in atmospheric precipitation, enter the soil and seas with them. Rain water on average contains about 0.00001% of bound N. It can be calculated that, in this way, up to 40 million tons of bound N. are introduced annually into the soil of the globe. 2. As early as Berthollet established that in the soil, without the introduction of nitrogenous compounds, the nitrogen content increases over time, due to the vital activity of certain species of bacteria. Subsequently these bacteria were isolated in pure cultures. These bacteria can enrich 1 hectare of soil with 48 kg of bound N. per year. "Besides these, free-living soil bacteria, the presence of symbiotically associated bacteria (Bacterium radicicola) was discovered in the nodular outgrowths on the roots of leguminous plants (Leguminosae), also capable of assimilating atmospheric N. and transferring this bound N. to the 'host plant' ('green' fertilizers—vetch, seradella, lupine, clover, etc.).- The introduction into the soil of nitrogen fertilizers in the form of Chilean saltpeter NaNO3, which undoubtedly has organic origin, or in the form of ammonium salts obtained during the coking of coal or in the production of illuminating gas from it—of course, cannot be considered a process enriching the resources of nature with respect to bound N., since these fertilizers are the heritage of dead animals and plants, and the presence of bound N. in them is the result of the fact that at one time plant or animal organisms assimilated compounds of bound N. in ready-made form. The reserves of Chilean saltpeter contain, according to sufficiently reliable calculations, about 30 million tons of bound N.; the world's coal reserves contain about 80 billion tons of bound N., but, with modern methods of extraction and utilization of compounds of bound N. from coal, only a small part of the reserves of bound N. contained in coal returns to the soil in the form of fertilizers: the larger part, when coal is burned in furnaces, undergoes decomposition processes with the formation of inert free N., thereby reducing the total reserve of bound N. in nature. The total reserve of bound N. in the soil of the globe can be estimated at approximately 40 billion tons. Finally, the chemical industry provides a very substantial replenishment of the reserves of bound N., returning (by technical fixation of free atmospheric nitrogen) to nature a fairly large amount of the bound nitrogen compounds taken from it. Opposite processes, changing the balance of the cycle of bound N. in nature, are processes in which compounds of bound N. are destroyed with the reverse formation of free N. These include unaccountable processes of decomposition of nitrogenous compounds when nitrogen-containing substances are burned, processes associated with the action of certain nitrogenous explosives that decompose upon explosion to free N., and bacterial denitrification processes, associated with the release of free N."