NITROGEN (Nitrogenium, N)- a chemical element of group V of the periodic system of elements of D. I. Mendeleev, atom, number 7, atomic mass 14.0067. Discovered by D. Rutherford in 1772. The following nitrogen isotopes are known (Table).

In various compounds, nitrogen has a variable valency, which can be equal to - 3, +1, +2, +3, +4 and +5.

distribution in nature. The total nitrogen content in the earth's crust is about 0.016 wt. %. Its main mass is in the air in a free, molecular form - N 2. Dry air contains on average 78.09% by volume (or 75.6% by weight) of free nitrogen. In relatively small quantities, free nitrogen is in a dissolved state in the waters of the oceans. Nitrogen in the form of compounds with other elements (bound nitrogen) is part of all plant and animal organisms.

Life is inextricably linked with the properties of easily changing complex nitrogenous substances - proteins. The composition of proteins on average includes 15-17% nitrogen. When organisms die, their complex nitrogenous compounds in the process of the nitrogen cycle turn into more simple connections: ammonia, ammonium salts, nitrites and nitrates. All nitrogen compounds, both organic and inorganic, found in the soil are collectively referred to as soil nitrogen.

Obtaining nitrogen

In laboratories, pure nitrogen is usually obtained by heating a concentrated aqueous solution of ammonium nitrate or a solution of a mixture of ammonium chloride and sodium nitrite:

NH 4 Cl + NaNO 2 = N 2 + NaCl + 2H 2 O.

In the technique of nitrogen with an admixture of up to 3% argon, it is obtained by fractionated distillation of liquid air.

Nitrogen Properties

In the free state, nitrogen is a colorless, odorless and tasteless gas, consisting of diatomic molecules - N 2 . The weight of 1 liter of it at t ° 0 ° and a pressure of 760 mm Hg. Art. equal to 1.2506 g, t° kip - 195.8°, t° pl - 209.86°; density of liquid A. 0.808 (at t ° - 195.8 °), solid - 1.026 (at t ° - 255 °). In 1 ml of water at t ° 0 °, 20 ° and 38 ° and a partial pressure of nitrogen equal to 760 mm, 0.0235, 0.0154 and 0.0122 ml of nitrogen are dissolved, respectively.

The solubility of nitrogen in the blood is less; it is at t ° 38 ° 0.0110 ml A. At low partial pressures of nitrogen, its solubility in the blood is somewhat greater than in water.

V normal conditions nitrogen is physiologically inert, but when air is inhaled, compressed to 2-2.5 atm, a state occurs called nitrogen anesthesia, similar to alcohol intoxication. This phenomenon can take place during diving operations (see) at a depth of several tens of meters. To prevent the occurrence of such a state, artificial gas mixtures are sometimes used in which nitrogen is replaced by helium or some other inert gas. With a sharp and significant decrease in the partial pressure of nitrogen, its solubility in the blood and tissues is so reduced that part of it is released in the form of bubbles, which is one of the causes of decompression sickness observed in divers when they quickly rise to the surface and in pilots at high takeoff speeds aircraft into the upper atmosphere (see Decompression sickness).

Application of nitrogen

Free nitrogen as a chemically inactive gas is used in laboratory practice and technology in all cases where the presence of oxygen in the surrounding atmosphere is unacceptable or undesirable, for example, when conducting a biological experiment under anaerobic conditions, when pouring large amounts of flammable liquids (to prevent fires) and so on. The bulk of free nitrogen is used in industry for the synthesis of ammonia, calcium cyanamide and nitric acid, which are the starting materials for the production of nitrogen fertilizers, explosives, paints, varnishes, pharmaceuticals and other.

Nitrogen compounds

Free nitrogen at ordinary temperatures is chemically inert; at high temperatures, it combines with many elements.

Nitrogen forms a number of compounds with hydrogen, the main of which are the following:

3. Nitrous acid (HN 3) - a colorless liquid boiling at t ° 37 ° with a pungent odor. Explodes with great force when heated. In aqueous solutions, it is stable and exhibits the properties of a weak acid. Its salts - azides - are unstable and explode when heated or hit. Lead azide Pb (N 3) 2 is used as a detonator. Inhalation of HN3 vapors causes severe headache and irritation of the mucous membranes.

With oxygen, nitrogen forms five oxides.

1. Nitrous oxide, or laughing gas (N 2 O), is a colorless gas, obtained by heating (above 190 °) ammonium nitrate:

NH 4 NO 3 \u003d N 2 O + 2H 2 O. In a mixture with oxygen, nitrous oxide is used as a weak drug that causes a state of intoxication, euphoria, dulling pain sensitivity. It is used for inhalation anesthesia (see).

2. Nitric oxide (NO) - a colorless gas, poorly soluble in water; in laboratories it is obtained by the action of nitric acid of medium concentration on copper:

8HNO 3 + 3Cu \u003d 2NO + 3Cu (NO 3) 2 + 4H 2 O, in technology - by blowing air through an electric arc flame. In air, it instantly oxidizes, forming red-brown vapors of nitrogen dioxide; together with the latter causes poisoning of the body (see below - Occupational hazards of nitrogen compounds).

3. Nitrogen dioxide (NO 2) - a red-brown gas that has a characteristic odor and consists of A. dioxide itself and its colorless polymer - nitrogen tetroxide (N 2 O 4) - nitrous anhydride. Nitrogen dioxide easily condenses into a red-brown liquid, boiling at t° 22.4° and solidifying at t° - 11° into colorless crystals. It dissolves in water to form nitrous and nitric acids:

2NO 2 + H 2 O \u003d HNO 2 + HNO 3.

It is a strong oxidizing agent and a dangerous poison. Nitrogen dioxide is formed during the production of nitric acid, during nitration reactions, pickling of metals, and the like, and therefore is an occupational poison.

4. Nitrogen trioxide, nitrous anhydride (N 2 O 3), is a dark blue liquid that hardens at t ° - 103 ° into blue crystals. Stable only at low temperatures. With water it forms a weak and fragile nitrous acid, with alkalis - salts of nitrous acid - nitrites.

5. Nitrogen pentoxide, nitric anhydride to-you (N 2 O 5), - colorless prismatic crystals having a density of 1.63, melting at t ° 30 ° into a yellow, slightly decomposing liquid; decomposition is enhanced by heating and exposure to light. The boiling point is about 50°. With water forms a strong, fairly stable nitric acid, with alkalis - salts of this acid - nitrates.

When heated, nitrogen combines directly with many metals, forming metal nitrides, for example, Li3N, Mg 3 N 2, AlN, etc. Many of them decompose with water to form ammonia, for example

Mg 3 N 2 + 6H 2 O \u003d 2NH 3 + 3Mg (OH) 2.

Nitrogen is part of a large number of organic compounds, among which alkaloids, amino acids, amines, nitro compounds, cyanide compounds and the most complex natural compounds - proteins are of particular importance.

Atmospheric nitrogen fixation. For a long time, natural Chilean nitrate and ammonia obtained by dry distillation served as starting materials for obtaining various nitrogen compounds necessary for agriculture, industry and military affairs. hard coal. With the depletion of the deposits of Chilean saltpeter, "nitrogen starvation" threatened humanity. The problem of nitrogen starvation was resolved in the late 19th and early 20th century by the development of a number of industrial methods for fixing atmospheric nitrogen. The most important of them is the synthesis of ammonia according to the scheme:

Determination of nitrogen

To determine free nitrogen, the gas to be analyzed is brought into contact with heated magnesium; in the presence of nitrogen, magnesium nitride is formed, which with water gives ammonia.

nitrogen cycle

Nitrogen is the most important biogenic element necessary for the construction of proteins and nucleic acids. However, atmospheric nitrogen is not available to animals and most plants. Therefore, in the nitrogen cycle, the process of its biological fixation (fixation of atmospheric molecular nitrogen) is of paramount importance. Nitrogen fixation is carried out by nitrogen-fixing microorganisms, such as bacteria from the genus Rhizobium, or nodule bacteria living in symbiosis (see) with legumes (peas, alfalfa, soybeans, lupins and others), on the roots of which nodules are formed containing bacteria that can absorb molecular nitrogen . Symbiotic nitrogen fixers also include some actinomycetes living in the root nodules of alder, sucker, sea buckthorn, and so on. Active nitrogen fixers are also some free-living microorganisms living in the soil, fresh and salt water bodies. This is an anaerobic spore-bearing bacterium Clostridium (Clostridium pasteurianum), discovered by S. N. Vinogradsky, an aerobic bacterium - Azotobacter (see Azotobacter). In addition, mycobacteria, some types of blue-green algae (Nostoc, Anabaena, etc.), as well as photosynthetic bacteria, have the ability to assimilate molecular nitrogen.

Nodule bacteria are of the greatest importance in soil enrichment with nitrogen. As a result of the activity of these bacteria, 100-250 kg/ha per season is introduced into the soil; blue-green algae in rice fields fix up to 200 kg/ha of nitrogen per year. Free-living nitrogen-fixing bacteria bind several tens of kilograms of nitrogen per hectare of soil.

S. N. Vinogradsky for the first time (1894) suggested that the initial product of the process of biological nitrogen fixation is ammonia. This assumption has now been fully confirmed. It has been proven that the conversion of N 2 to NH 3 is an enzymatic process. The enzyme that carries out this process (nitrogenase) consists of two protein components, is active only in the absence of oxygen, and the process itself occurs due to the energy of adenosine triphosphoric acid (ATP). Plants, as well as microorganisms, then convert inorganic ammonium nitrogen into its organic compounds (amino acids, proteins, nucleic acids, and so on), and in this form it becomes available to animals and humans, being included in the metabolic processes occurring in their organisms. Organic nitrogen of animals and plants enters the soil (with animal secretions or their decomposition products) and is processed by various worms, mollusks, nematodes, insects, and microorganisms living there. Soil microorganisms - ammonifiers (putrefactive bacteria, some actinomycetes and fungi) - mineralize in turn organic nitrogen soil (the bodies of animals and plants, organic fertilizers, humus) to ammonium. Ammonification is a complex of enzymatic processes occurring mainly in two stages: the hydrolysis of proteins and nucleic acids to amino acids and nitrogenous bases and the subsequent decomposition of these compounds to ammonia. The resulting ammonia is neutralized by reacting with the organic and inorganic acids contained in the soil. This results in the formation of ammonium salts. Ammonium salts and ammonia, in turn, undergo nitrification under the influence of nitrifying bacteria (discovered in 1890 by S. N. Vinogradsky) with the formation of nitrates and nitrites.

The processes of nitrification and ammonification provide plants with easily digestible nitrogen compounds. Ammonium salts and nitrates are absorbed by plants and microorganisms, turning into nitrogen organic compounds. However, part of the nitrogen is converted in the soil into molecular nitrogen as a result of the denitrification process carried out by microorganisms living in the soil - denitrifiers (Fig.). Denitrifying bacteria are widely distributed in nature, found in large numbers in soil, manure and in smaller numbers in the water of rivers, lakes and seas. The most typical denitrifiers are mobile, Gram-negative rods. These include Bacterium fluorescens, B. denitrificans, B. pyocyaneum and more.

The process of denitrification leads to the loss of nitrogen available to plants, however, the constantly ongoing process of nitrogen fixation to some extent compensates for these losses, and under certain conditions (in particular, when the soil is rich in nitrogen-free organic substances), it significantly enriches the soil with bound nitrogen.

In general, the combined effect of the processes of nitrogen fixation, nitrification and denitrification is of great biogeochemical importance, contributing to the maintenance of a dynamic balance between the content of molecular nitrogen in the atmosphere and the bound nitrogen of the soil, flora and fauna.

The nitrogen cycle thus plays a critical role in sustaining life on Earth.

Occupational hazards of nitrogen compounds

Nitric acid (see), ammonia (see), amino compounds (see Amines) and amido compounds (see Amides), as well as mixtures of nitrogen oxides, or nitrogases (N 2 O, NO, NO 2 , N 2 O 4 and N 2 O 5). The latter are formed during the production and use of nitric acid (in the process of its interaction with various metals or organic substances), in the process of thermal oxidation of air nitrogen during electric and gas welding, the operation of diesel and carburetor engines, fuel combustion in powerful boiler houses, as well as during blasting. etc. The general nature of the action of nitrogases on the body depends on the content of various nitrogen oxides in the gas mixture. Basically, poisoning proceeds by an irritating, or nitrite, type of action. When nitrogen oxides come into contact with the moist surface of the lungs, nitric and nitrous acids are formed, which affect the lung tissue, causing pulmonary edema. At the same time, nitrates (see) and nitrites (see) are formed in the blood, directly acting on the blood vessels, expanding them and causing a decrease in blood pressure. Nitrites, interacting with oxyhemoglobin, turn it into methemoglobin, causing methemoglobinemia (see). A common consequence of the action of nitrogen oxides is oxygen deficiency.

Under production conditions, there may be cases of exposure to individual oxides of nitrogen (see below).

Nitrous oxide. Its large concentrations cause tinnitus, asphyxia, loss of consciousness. Death occurs from paralysis of the respiratory center.

Nitric oxide acts on the central nervous system, affects hemoglobin (converts oxyhemoglobin to methemoglobin).

With mild nitric oxide poisoning, general weakness, drowsiness, dizziness are observed (symptoms are reversible).

With more severe poisoning, the initial symptoms intensify, they are joined by nausea, sometimes vomiting, and a fainting state occurs. With moderate poisoning, severe weakness and dizziness last for many hours, cyanosis of the mucous membranes and skin, and increased heart rate are often observed. In severe poisoning, the initial symptoms often subside, but after a 1-3-day remission, weakness and dizziness appear, a decrease in blood pressure, a gray-blue color of the mucous membranes and skin, an increase and soreness of the liver are observed; the borders of the heart are expanded, the tones are deaf, the pulse is slow. There are polyneuritis, polyneuralgia. Chocolate-brown blood, high viscosity. The consequences of severe poisoning can last more than a year: impaired associative abilities, weakening of memory and muscle strength, general weakness, headache, dizziness, fatigue.

Nitrogen dioxide. Acute poisoning begins with a mild cough, severe cases- with a strong cough, chest tightness, headache, sometimes vomiting, salivation. The period of relatively satisfactory condition lasts 2-18 hours. Then there are signs of increasing pulmonary edema: severe weakness, increasing cough, chest pain, cyanosis, many wet rales in the lungs, rapid heartbeat, sometimes chills, fever. Significant disorders of the gastrointestinal tract are not uncommon: nausea, vomiting, diarrhea, severe pain in the upper abdomen. Pulmonary edema is characterized by a serious condition (severe cyanosis, severe shortness of breath, rapid pulse, cough with frothy sputum, sometimes with blood). Blood pressure is normal, in the blood - an increase in the number of erythrocytes and hemoglobin, leukocytosis, delayed ESR. X-ray - reduced transparency of the lung fields, in both lungs a large number of flocculent opacities of various sizes. Toxic pulmonary edema is accompanied by a "blue" type of hypoxemia, with a complication of collapse, a "gray" type is observed (see Hypoxia). Frequent complications of pneumonia. Possible fatal outcome. On the section - pulmonary edema, hemorrhages in them, dark liquid blood in the heart and blood vessels. The condition of the poisoned and the prognosis worsens if the victims suffered from heart or lung diseases before the poisoning.

In chronic poisoning - chronic inflammatory diseases of the upper respiratory tract, chronic bronchitis, emphysema, lowering blood pressure, greenish plaque on the teeth, destruction of the crowns of the incisors.

Nitrous anhydride acts on the body in a similar way to nitric oxide and its other lower oxides.

First aid for poisoning with nitrogen compounds- move the victim to fresh air; ensure complete rest, inhalation of oxygen. According to indications - cardiac drugs, when breathing stops - lobelin. Then the mandatory transportation of the victim in the supine position to the hospital. With signs of incipient pulmonary edema - intravenously 10-20 ml of a 10% solution of calcium chloride, 20 ml of a 40% glucose solution with ascorbic acid (500 mg), oxygen therapy.

Treatment of developed pulmonary edema depends on the type of hypoxemia. With the "blue" type - intermittent administration of oxygen (carbogen is contraindicated), bloodletting (200-300 ml), if necessary - repeating it after 6-8 hours; blood pressure lowering agents, cardiac agents are recommended. With the "gray" type of anoxemia - stimulation of the respiratory and vasomotor center by intermittent inhalation of carbogen, caffeine, ephedrine, intravenously 50-100 ml of 40% glucose solution. Bloodletting is contraindicated.

In order to prevent and treat pneumonia - early appointment of sulfonamides and antibiotics.

Prevention: personal protection - filtering gas masks of grades B, M, KB, acid-proof gloves and boots, sealed goggles, special clothing. There is a need for complete sealing of production equipment where nitrogases can be formed and released, shelter for fixed sources of these gases, and a local ventilation system.

The maximum permissible concentration for nitrogen oxides in the air of working premises is 5 mg / m 3 (in terms of NO 2), in the atmospheric air of settlements 0.085 mg / m 3 or 0.4 mg / m 3 (for nitric acid).

The determination of nitrogen oxides in the air is based on the absorption of nitrogen dioxide and nitrogen tetroxide by a solution of potassium iodide and the colorimetric determination of the formed nitrous acid with the Griess-Iloshvai reagent.

Bibliography: Nekrasov B.V. Fundamentals of General Chemistry, t. 1, p. 377, M., 1969; Remy G. Course of inorganic chemistry, trans. from German, vol. 1, p. 560, M., 1972.

The circle A.- Vinogradsky S. N. Soil microbiology, M., 1952; Kretovich V. L. Exchange of nitrogen in plants, M., 1972, bibliogr.; Mishustin E. N. and Shilnikova V. K. Biological fixation of atmospheric nitrogen, M., 1968, bibliogr.

Occupational hazards of compounds A. - Harmful substances in industry, ed. N. V. Lazareva, part 2, p. 136, L., 1971; Occupational health in the chemical industry, ed. Z. A. Volkova and others, p. 373, M., 1967; Yu. A. Gurtovoy. Poisoning with vapors of nitric acid, Sud.-med. examination, vol. 12, no. 3, p. 45, 1969; Neimark E. Z. and Singer F. X. Occupational poisoning of coal mine workers, their treatment and prevention, p. 34, Moscow, 1961; Peregud E. A., Bykhovskaya M. S. and Gernet E. V. Rapid methods for determining harmful substances in the air, p. 67, M., 1970; Safronov V. A. Features of the clinical course of pulmonary edema in combined lesions with nitric acid, Voyen.-med. journal, no. 7, p. 32, 1966; Air quality criteria for nitrogen oxides, Washington, 1971, bibliogr.

V. P. Mishin; Z. G. Evstigneeva, V. L. Kretovich (circulation of A.); E. N. Marchenko (prof.).

The content of the article

NITROGEN, N (nitrogenium), chemical element (at. number 7) VA subgroup of the Periodic Table of Elements. The Earth's atmosphere contains 78% (vol.) nitrogen. To show how large these reserves of nitrogen are, we note that there is so much nitrogen in the atmosphere above each square kilometer of the earth's surface that up to 50 million tons of sodium nitrate or 10 million tons of ammonia (a combination of nitrogen with hydrogen) can be obtained from it, and yet this is a small fraction of the nitrogen contained in the earth's crust. The existence of free nitrogen indicates its inertness and the difficulty of interacting with other elements at ordinary temperatures. Bound nitrogen is part of both organic and inorganic matter. Plant and animal life contains nitrogen bound to carbon and oxygen in proteins. In addition, nitrogen-containing inorganic compounds are known and can be obtained in large quantities, such as nitrates (NO 3 -), nitrites (NO 2 -), cyanides (CN -), nitrides (N 3 -) and azides (N 3 - ).

History reference.

The experiments of A. Lavoisier, devoted to the study of the role of the atmosphere in maintaining life and combustion processes, confirmed the existence of a relatively inert substance in the atmosphere. Not having established the elemental nature of the gas remaining after combustion, Lavoisier called it azote, which in ancient Greek means "lifeless". In 1772, D. Rutherford from Edinburgh established that this gas is an element and called it "harmful air". The Latin name for nitrogen comes from the Greek words nitron and gen, which means "forming saltpetre".

Nitrogen fixation and the nitrogen cycle.

The term "nitrogen fixation" refers to the process of fixing atmospheric nitrogen N 2 . In nature, this can happen in two ways: either legumes, such as peas, clover, and soybeans, accumulate nodules on their roots, in which nitrogen-fixing bacteria convert it into nitrates, or atmospheric nitrogen is oxidized by oxygen under conditions of a lightning discharge. S. Arrhenius found that up to 400 million tons of nitrogen are fixed in this way annually. In the atmosphere, nitrogen oxides combine with rainwater to form nitric and nitrous acids. In addition, it has been established that with rain and snow, approx. 6700 g of nitrogen; reaching the soil, they turn into nitrites and nitrates. Plants use nitrates to form plant proteins. Animals, eating these plants, assimilate the protein substances of plants and turn them into animal proteins. After the death of animals and plants, they decompose, nitrogen compounds turn into ammonia. Ammonia is used in two ways: bacteria that do not form nitrates break it down to elements, releasing nitrogen and hydrogen, and other bacteria form nitrites from it, which are oxidized to nitrates by other bacteria. Thus, the nitrogen cycle in nature, or the nitrogen cycle, occurs.

The structure of the nucleus and electron shells.

In nature, there are two stable nitrogen isotopes: with a mass number of 14 (contains 7 protons and 7 neutrons) and with a mass number of 15 (contains 7 protons and 8 neutrons). Their ratio is 99.635:0.365, so the atomic mass of nitrogen is 14.008. Unstable nitrogen isotopes 12 N, 13 N, 16 N, 17 N were obtained artificially. Schematically, the electronic structure of the nitrogen atom is as follows: 1 s 2 2s 2 2p x 1 2py 1 2pz one . Therefore, on the outer (second) electron shell there are 5 electrons that can participate in the formation chemical bonds; nitrogen orbitals can also accept electrons, i.e. the formation of compounds with oxidation states from (–III) to (V) is possible, and they are known.

Molecular nitrogen.

From the definitions of gas density, it was established that the nitrogen molecule is diatomic, i.e. the molecular formula of nitrogen is Nє N (or N 2). Two nitrogen atoms have three external 2 p-electron of each atom form a triple bond: N::: N:, forming electron pairs. The measured N–N interatomic distance is 1.095 Å. As in the case of hydrogen ( cm. HYDROGEN) , there are nitrogen molecules with different nuclear spins - symmetric and antisymmetric. At ordinary temperature, the ratio of symmetric and antisymmetric forms is 2:1. In the solid state, two modifications of nitrogen are known: a- cubic and b– hexagonal with transition temperature a ® b-237.39 ° C. Modification b melts at -209.96°C and boils at -195.78°C at 1 atm ( cm. tab. one).

The dissociation energy of a mole (28.016 g or 6.023 H 10 23 molecules) of molecular nitrogen into atoms (N 2 2N) is approximately –225 kcal. Therefore, atomic nitrogen can be formed in a quiet electrical discharge and is chemically more active than molecular nitrogen.

Receipt and application.

The method of obtaining elemental nitrogen depends on the required purity. Huge amounts of nitrogen are obtained for the synthesis of ammonia, while small admixtures of noble gases are acceptable.

nitrogen from the atmosphere.

Economically, the release of nitrogen from the atmosphere is due to the cheapness of the method of liquefying purified air (water vapor, CO 2 , dust, other impurities removed). Successive cycles of compression, cooling and expansion of such air lead to its liquefaction. Liquid air is subjected to fractional distillation with a slow rise in temperature. Noble gases are released first, then nitrogen, and liquid oxygen remains. Purification is achieved by multiple fractionation processes. This method produces many millions of tons of nitrogen annually, mainly for the synthesis of ammonia, which is the feedstock in the technology for the production of various nitrogen-containing compounds for industry and agriculture. In addition, a purified nitrogen atmosphere is often used when the presence of oxygen is unacceptable.

laboratory methods.

Small amounts of nitrogen can be obtained in the laboratory different ways, oxidizing ammonia or an ammonium ion, for example:

The process of oxidation of the ammonium ion with the nitrite ion is very convenient:

Other methods are also known - the decomposition of azides when heated, the decomposition of ammonia with copper (II) oxide, the interaction of nitrites with sulfamic acid or urea:

With the catalytic decomposition of ammonia at high temperatures, nitrogen can also be obtained:

physical properties.

Some physical properties and nitrogen are given in table. one.

Table 1. SOME PHYSICAL PROPERTIES OF NITROGEN
Density, g / cm 3	0.808 (liquid)
Melting point, °С	–209,96
Boiling point, °C	–195,8
Critical temperature, °C	–147,1
Critical pressure, atm a	33,5
Critical density, g / cm 3 a	0,311
Specific heat capacity, J/(molChK)	14.56 (15°C)
Electronegativity according to Pauling	3
covalent radius,	0,74
crystal radius,	1.4 (M 3–)
Ionization potential, V b
first	14,54
second	29,60
a The temperature and pressure at which the densities of liquid and gaseous nitrogen are the same. b The amount of energy required to remove the first outer and following electrons, per 1 mole of atomic nitrogen.

Chemical properties.

As already noted, the predominant property of nitrogen under normal conditions of temperature and pressure is its inertness, or low chemical activity. The electronic structure of nitrogen contains an electron pair for 2 s-level and three half-filled 2 R-orbitals, so one nitrogen atom can bind no more than four other atoms, i.e. its coordination number is four. The small size of an atom also limits the number of atoms or groups of atoms that can be bonded to it. Therefore, many compounds of other members of the VA subgroup either have no analogues among nitrogen compounds at all, or similar nitrogen compounds turn out to be unstable. Thus, PCl 5 is a stable compound, while NCl 5 does not exist. The nitrogen atom is able to bond with another nitrogen atom, forming several fairly stable compounds, such as hydrazine N 2 H 4 and metal azides MN 3 . This type of bond is unusual for chemical elements (with the exception of carbon and silicon). At elevated temperatures, nitrogen reacts with many metals to form partially ionic nitrides M x N y. In these compounds, nitrogen is negatively charged. In table. 2 shows the oxidation states and examples of the corresponding compounds.

Nitrides.

Nitrogen compounds with more electropositive elements, metals and non-metals - nitrides - are similar to carbides and hydrides. They can be divided depending on the nature of the M–N bond into ionic, covalent, and with an intermediate type of bond. As a rule, these are crystalline substances.

Ionic nitrides.

The bond in these compounds involves the transfer of electrons from the metal to nitrogen with the formation of the N 3– ion. These nitrides include Li 3 N, Mg 3 N 2 , Zn 3 N 2 and Cu 3 N 2 . In addition to lithium, other alkali metals of the IA subgroup do not form nitrides. Ionic nitrides have high melting points, react with water to form NH 3 and metal hydroxides.

covalent nitrides.

When the electrons of nitrogen participate in the formation of a bond together with the electrons of another element without transferring them from nitrogen to another atom, nitrides with a covalent bond are formed. Hydrogen nitrides (eg ammonia and hydrazine) are fully covalent, as are nitrogen halides (NF 3 and NCl 3). Covalent nitrides include, for example, Si 3 N 4 , P 3 N 5 and BN - highly stable white substances, and BN has two allotropic modifications: hexagonal and diamond-like. The latter is formed at high pressures and temperatures and has a hardness close to that of diamond.

Nitrides with an intermediate type of bond.

Transition elements react with NH 3 at high temperature to form an unusual class of compounds in which nitrogen atoms are distributed between regularly spaced metal atoms. There is no clear displacement of electrons in these compounds. Examples of such nitrides are Fe 4 N, W 2 N, Mo 2 N, Mn 3 N 2. These compounds are generally completely inert and have good electrical conductivity.

Hydrogen compounds of nitrogen.

Nitrogen and hydrogen interact to form compounds that vaguely resemble hydrocarbons. The stability of hydrogen nitrogens decreases with an increase in the number of nitrogen atoms in the chain, in contrast to hydrocarbons, which are also stable in long chains. The most important hydrogen nitrides are ammonia NH 3 and hydrazine N 2 H 4 . These also include hydronitrous acid HNNN (HN 3).

Ammonia NH3.

Ammonia is one of the most important industrial products of the modern economy. At the end of the 20th century The US produced approx. 13 million tons of ammonia annually (in terms of anhydrous ammonia).

The structure of the molecule.

The NH 3 molecule has an almost pyramidal structure. The H–N–H bond angle is 107°, which is close to the tetrahedral angle of 109°. The unshared electron pair is equivalent to the attached group, as a result, the coordination number of nitrogen is 4 and nitrogen is located in the center of the tetrahedron.

properties of ammonia.

Some physical properties of ammonia in comparison with water are given in table. 3.

The boiling and melting points of ammonia are much lower than those of water, despite the similarity of molecular weights and the similarity of the structure of the molecules. This is due to the relatively greater strength of intermolecular bonds in water than in ammonia (such an intermolecular bond is called hydrogen).

ammonia as a solvent.

The high dielectric constant and dipole moment of liquid ammonia make it possible to use it as a solvent for polar or ionic non organic matter. Ammonia solvent occupies an intermediate position between water and organic solvents such as ethyl alcohol. Alkali and alkaline earth metals dissolve in ammonia, forming dark blue solutions. It can be assumed that solvation and ionization of valence electrons occurs in solution according to the scheme

The blue color is associated with solvation and the movement of electrons or with the mobility of "holes" in a liquid. At a high concentration of sodium in liquid ammonia, the solution takes on a bronze color and is characterized by high electrical conductivity. The unbound alkali metal can be separated from such a solution by evaporation of the ammonia or by adding sodium chloride. Solutions of metals in ammonia are good reducing agents. Autoionization occurs in liquid ammonia

similar to the process taking place in water:

Some Chemical properties both systems are compared in table. 4.

Liquid ammonia as a solvent has an advantage in some cases where it is impossible to carry out reactions in water due to the rapid interaction of components with water (for example, oxidation and reduction). For example, in liquid ammonia, calcium reacts with KCl to form CaCl 2 and K, since CaCl 2 is insoluble in liquid ammonia, but K is soluble, and the reaction proceeds completely. In water, such a reaction is impossible due to the rapid interaction of Ca with water.

Getting ammonia.

Gaseous NH 3 is released from ammonium salts under the action of a strong base, for example, NaOH:

The method is applicable in laboratory conditions. Small ammonia production is also based on the hydrolysis of nitrides, such as Mg 3 N 2 , with water. Calcium cyanamide CaCN 2 also forms ammonia when interacting with water. The main industrial method for producing ammonia is its catalytic synthesis from atmospheric nitrogen and hydrogen at high temperature and pressure:

Hydrogen for this synthesis is obtained by thermal cracking of hydrocarbons, the action of water vapor on coal or iron, the decomposition of alcohols with water vapor, or the electrolysis of water. Many patents have been obtained for the synthesis of ammonia, differing in the process conditions (temperature, pressure, catalyst). There is a method of industrial production during the thermal distillation of coal. The names of F. Haber and K. Bosch are associated with the technological development of ammonia synthesis.

Table 4. COMPARISON OF REACTIONS IN WATER AND AMMONIA MEDIUM
Water environment	Ammonia medium
Neutralization
OH - + H 3 O + ® 2H 2 O	NH 2 - + NH 4 + ® 2NH 3
Hydrolysis (protolysis)
PCl 5 + 3H 2 O POCl 3 + 2H 3 O + + 2Cl –	PCl 5 + 4NH 3 PNCl 2 + 3NH 4 + + 3Cl -
substitution
Zn + 2H 3 O + ® Zn 2+ + 2H 2 O + H 2	Zn + 2NH 4 + ® Zn 2+ + 2NH 3 + H 2
solvation (complexation)
Al 2 Cl 6 + 12H 2 O 2 3+ + 6Cl -	Al 2 Cl 6 + 12NH 3 2 3+ + 6Cl -
Amphoteric
Zn 2+ + 2OH - Zn (OH) 2	Zn 2+ + 2NH 2 - Zn (NH 2) 2
Zn(OH) 2 + 2H 3 O + Zn 2+ + 4H 2 O	Zn(NH 2) 2 + 2NH 4 + Zn 2+ + 4NH 3
Zn(OH) 2 + 2OH – Zn(OH) 4 2–	Zn(NH 2) 2 + 2NH 2 – Zn(NH 2) 4 2–

Chemical properties of ammonia.

In addition to the reactions mentioned in Table. 4, ammonia reacts with water, forming the compound NH 3 H H 2 O, which is often mistakenly considered ammonium hydroxide NH 4 OH; in fact, the existence of NH 4 OH in solution has not been proven. An aqueous solution of ammonia (" ammonia”) consists mainly of NH 3, H 2 O and small concentrations of NH 4 + and OH - ions formed during dissociation

The main nature of ammonia is explained by the presence of a lone electron pair of nitrogen: NH 3 . Therefore, NH 3 is a Lewis base, which has the highest nucleophilic activity, manifested in the form of association with a proton, or the nucleus of a hydrogen atom:

Any ion or molecule capable of accepting an electron pair (an electrophilic compound) will react with NH 3 to form a coordination compound. For instance:

Symbol M n+ represents a transition metal ion (B-subgroups of the periodic table, for example, Cu 2+ , Mn 2+ , etc.). Any protic (i.e. H-containing) acid reacts with ammonia in aqueous solution to form ammonium salts, such as ammonium nitrate NH 4 NO 3, ammonium chloride NH 4 Cl, ammonium sulfate (NH 4) 2 SO 4, phosphate ammonium (NH 4) 3 PO 4. These salts are widely used in agriculture as fertilizers to introduce nitrogen into the soil. Ammonium nitrate is also used as an inexpensive explosive; for the first time it was applied with fuel oil (diesel oil). An aqueous solution of ammonia is used directly for introduction into the soil or with irrigation water. Urea NH 2 CONH 2 obtained by synthesis from ammonia and carbon dioxide, is also a fertilizer. Gaseous ammonia reacts with metals such as Na and K to form amides:

Ammonia reacts with hydrides and nitrides also to form amides:

Alkali metal amides (for example, NaNH 2) react with N 2 O when heated, forming azides:

Gaseous NH 3 reduces heavy metal oxides to metals at high temperature, apparently due to the hydrogen formed as a result of the decomposition of ammonia into N 2 and H 2:

The hydrogen atoms in the NH 3 molecule can be replaced by a halogen. Iodine reacts with a concentrated solution of NH 3 , forming a mixture of substances containing NI 3 . This substance is very unstable and explodes at the slightest mechanical impact. The reaction of NH 3 with Cl 2 produces chloramines NCl 3 , NHCl 2 and NH 2 Cl. When exposed to ammonia sodium hypochlorite NaOCl (formed from NaOH and Cl 2), the end product is hydrazine:

Hydrazine.

The above reactions are a method for obtaining hydrazine monohydrate of the composition N 2 H 4 H H 2 O. Anhydrous hydrazine is formed by special distillation of the monohydrate with BaO or other water-removing substances. In terms of properties, hydrazine slightly resembles hydrogen peroxide H 2 O 2. Pure anhydrous hydrazine is a colorless hygroscopic liquid boiling at 113.5°C; dissolves well in water, forming a weak base

In an acidic environment (H +), hydrazine forms soluble hydrazonium salts of the + X - type. The ease with which hydrazine and some of its derivatives (eg, methylhydrazine) react with oxygen allows it to be used as a component of liquid propellant. Hydrazine and all its derivatives are highly toxic.

nitrogen oxides.

In compounds with oxygen, nitrogen exhibits all oxidation states, forming oxides: N 2 O, NO, N 2 O 3, NO 2 (N 2 O 4), N 2 O 5. There is little information on the formation of nitrogen peroxides (NO 3 , NO 4). 2HNO2. Pure N 2 O 3 can be obtained as a blue liquid at low temperatures (–20

At room temperature, NO 2 is a dark brown gas that has magnetic properties due to the presence of an unpaired electron. At temperatures below 0 ° C, the NO 2 molecule dimerizes into dinitrogen tetroxide, and at –9.3 ° C, the dimerization proceeds completely: 2NO 2 N 2 O 4. In the liquid state, only 1% NO 2 is not dimerized, and at 100° C 10% N 2 O 4 remains in the form of a dimer.

NO 2 (or N 2 O 4) reacts in warm water with the formation of nitric acid: 3NO 2 + H 2 O \u003d 2HNO 3 + NO. The NO 2 technology is therefore very important as an intermediate step in the production of an industrially important product, nitric acid.

Nitric oxide(V)

N 2 O 5 ( outdated. nitric anhydride) - white crystalline substance, obtained by dehydration of nitric acid in the presence of phosphorus oxide P 4 O 10:

2MX + H 2 N 2 O 2 . When the solution is evaporated, a white explosive is formed with the proposed structure H–O–N=N–O–H.

Nitrous acid

HNO 2 does not exist in its pure form, however, aqueous solutions of its low concentration are formed by adding sulfuric acid to barium nitrite:

Nitrous acid is also formed by dissolving an equimolar mixture of NO and NO 2 (or N 2 O 3) in water. Nitrous acid is slightly stronger than acetic acid. The degree of nitrogen oxidation in it is +3 (its structure is H–O–N=O), i.e. it can be both an oxidizing agent and a reducing agent. Under the action of reducing agents, it is usually reduced to NO, and when interacting with oxidizing agents, it is oxidized to nitric acid.

The rate of dissolution of certain substances, such as metals or iodide ion, in nitric acid depends on the concentration of nitrous acid present as an impurity. Salts of nitrous acid - nitrites - dissolve well in water, except for silver nitrite. NaNO 2 is used in the production of dyes.

Nitric acid

HNO 3 is one of the most important inorganic products in the mainstream chemical industry. It is used in the technology of many other inorganic and organic substances, such as explosives, fertilizers, polymers and fibers, dyes, pharmaceuticals, etc.

Literature:

Azotchik's Handbook. M., 1969
Nekrasov B.V. Fundamentals of General Chemistry. M., 1973
Problems of nitrogen fixation. Inorganic and physical chemistry . M., 1982



NITROGEN, N (French Az), a chemical element (Nitrogenium - from nitrum, saltpeter, "forming saltpeter"; in German - Stickstoff "suffocating gas", in French - Azote, from Greek α - negation, ξωη - life , lifeless); atomic weight 14.009, serial number 7.

Physical properties. D of pure nitrogen (at D of air = 1) 0.9674; but usually we are dealing with nitrogen from the air, with a content of 1.12% argon, D of such nitrogen is 0.9721; the weight of 1 liter of pure nitrogen at 0°C and 760 mm is 1.2507 g, the weight of 1 liter of "atmospheric" nitrogen is 1.2567 g. The solubility of nitrogen in water is less than the solubility of oxygen. 1 liter of water at 760 mm and 0 ° C dissolves 23.5 cm 3 of nitrogen (O 2 solubility - 48.9 cm 3), at 20 ° C - 15.4 cm 3 of nitrogen (O 2 solubility - 31.0 cm 3 ). Charcoal freshly calcined absorbs, according to Dewar, in 1 cm 3 at 0 ° C only 15 cm 3 of nitrogen, at -185 ° C it absorbs 155 cm 3 of nitrogen (volumes are listed at 0 ° C and 760 mm). The critical temperature is -147 ° C at a critical pressure of 33 atm., or 25 m of mercury, the boiling point at 760 mm is -195 °.67 ± 0 °.05, and the melting point at 88 mm ± 4 mm is - 210 ° .52±0°.2. The expansion coefficient of nitrogen at 1 atm is 0.003667; specific heat at 20°C is 0.249, and for the temperature range (0-1400)°C, on average, 0.262; ratio with p /c η = 1.40, as for O 2 . Liquid nitrogen is colorless, mobile like water, although lighter than the latter. The specific gravity at the boiling point and 760 mm is 0.7914, at -184°C - 0.7576, at -195.5°C - 0.8103 and at -205°C - 0.8537; near the freezing point - 0.8792 (figures fluctuate depending on the Ar content). Specific heat of liquid nitrogen between -196°C and -208°C - 0.430; the heat of vaporization of 1 kg of liquid nitrogen at a boiling point of -195°.55 is 47.65 Cal. From 1 liter of liquid nitrogen during evaporation, at atmospheric pressure and 0°C, 14°C and 27°C, respectively: 640, 670 and 700 liters of gaseous nitrogen are formed. Liquid nitrogen is non-magnetic and does not conduct electricity.

Chemical properties nitrogen is largely determined by its extreme inertness under ordinary conditions of temperature and pressure, due to the stability of N 2 molecules. Only lithium metal combines with nitrogen at a low temperature, releasing 69000 cal and forming lithium nitride NLi 3 . Nitride Ba is formed at 560°C and has the formula Ba 3 N 2 ; about other nitrides. Both with oxygen and with hydrogen, nitrogen combines only at high temperatures, and the reaction with oxygen is endothermic, and with hydrogen it is exothermic. The valence of nitrogen is determined by the structure of its atom according to Bohr. When all five electrons are removed from the outer ring, nitrogen becomes a five-charged positive ion; when the upper ring is replenished with three electrons up to the limiting number - eight - the nitrogen atom appears as a three-charged electronegative ion. The state of nitrogen in ammonium compounds can be easily elucidated by the theory of complex compounds. Nitrogen gives a whole series of compounds with oxygen and with halides (the latter compounds are extremely explosive due to the strong endothermicity of their formation). With hydrogen, nitrogen gives compounds: ammonia and hydrazoic acid. In addition, the following are known: the combination of nitrogen with hydrogen - hydrazine and with hydrogen and oxygen - hydroxylamine.

Application of nitrogen. Gaseous nitrogen is used as an inert gas in medicine for immobilizing areas of the lungs affected by tuberculosis (Pneumotorax operation), for protecting metals from the chemical action of active gases on them, and in general in cases where it is necessary to prevent any unwanted chemical reaction(for example, to fill incandescent light bulbs, to inflate automobile rubber tires that are attacked by air at high pressure, to preserve the colors of valuable paintings placed in hermetic vessels filled with nitrogen, to prevent a fire hazard when pouring gasoline and other flammable liquids, and etc.). But the most important technical application of nitrogen is in the process of obtaining synthetic ammonia from the elements.

When evaluating the properties of nitrogen and its exceptional importance in the general economy of organic nature and human social life, one should sharply distinguish between free nitrogen and bound nitrogen, that is, nitrogen that has already entered into a chemical combination with some other element, ch. arr. with oxygen, hydrogen and carbon. Free nitrogen, under the conditions of temperature and pressure prevailing on the surface of the globe, is an extremely inert element. The mouse in the classical experiment of Lavoisier died in oxygen-deprived air, that is, in almost pure nitrogen. Meanwhile, the bound nitrogen is, as it were, the carrier of life, for all living beings, without exception, be they plants or animals, build their organism necessarily with the participation of the so-called. protein substances, inevitably containing in their chemical composition nitrogen (proteins contain up to 16% nitrogen). The process of transition from free nitrogen to bound nitrogen and vice versa is a process of the greatest importance in nature and the greatest problem of agriculture, and, more recently, of industry. Free nitrogen is contained in a mixture with other gases in the atmosphere in an immense amount, accounting for about 4/5 by volume (75.51 weight%) of the entire atmosphere and enveloping the globe with air cover, gradually becoming more and more rarefied, reaching a height of tens of kilometers . Over one hectare of the earth's surface contains so much nitrogen that, if it were in a bound state, it would be enough to provide all living nature and the needs of mankind for 20 years (A. E. Moser). But free nitrogen can only with a huge effort. forced to combine with other elements, and moreover, not only in those cases when this combination occurs endothermally (as, for example, in the formation of oxygen compounds of nitrogen), but also in those cases when the combination of nitrogen with another element is accompanied by the release of energy and is a reaction exothermic (combination of nitrogen with hydrogen).

Only in exceptional cases, for example, with lithium, the combination of nitrogen proceeds easily under ordinary conditions of temperature and pressure. Therefore, in the general balance of bound nitrogen in nature, one has to state a cycle. Plants take up bound nitrogen in the form of soluble salts from the soil and make proteins; animals use ready-made nitrogen compounds during metabolism due to absorbed plant foods, releasing bound nitrogen compounds, unassimilated, and also formed as a result of the breakdown of protein substances in their body - in excrement and urine, and, finally, introducing their entire body upon their death into the total balance of bound nitrogen in nature for further processes mineralization of protein and other nitrogenous substances occurring in the soil. In these latter processes, an enormous role remains with soil microorganisms, as a result of whose vital activity complex nitrogenous organic compounds are converted into the simplest salts of nitric acid, which, in turn, is formed as a result of the oxidation of ammonia compounds in the soil as an earlier stage in the destruction of protein substances and id products. decay. Taking into account the extreme inertness of free nitrogen, which is unable to enter into compounds on its own, and, on the other hand, losses or cases of deep destruction of a nitrogenous compound to free nitrogen (for example, as a result of the vital activity denitrifying soil bacteria, when burning coal, firewood and peat, when nitrogenous compounds are washed out of the soil by rain into rivers and seas, when the garbage of large cities descends into rivers, etc.), one could consider the gradual impoverishment of nature as an inevitable consequence of all this bound nitrogen and, as a result, the death of organic life on earth, if some processes had not flowed into the general channel of the cycle of bound nitrogen, replenishing the indicated loss of bound nitrogen in nature. Such a natural source of bound nitrogen in nature is atmospheric precipitation, which brings nitrogen oxides into the soil, formed in the atmosphere during electrical discharges, which force a certain amount of atmospheric nitrogen to combine with oxygen ( rainwater contains about 0.00001% bound nitrogen). It can be calculated that up to 400 million tons of bound nitrogen are annually introduced into the soil of the globe in this way. In addition, Berthelot was able to establish that in the soil, without introducing new reserves of nitrogenous compounds into it, the nitrogen content increases over time due to the vital activity of certain types of bacteria. Subsequently, these bacteria were isolated in pure cultures, namely: the anaerobic bacterium of butyric fermentation (Clostridium pasteurianum) and the aerobic bacterium (Azotobakter Winogradsky, which can enrich the soil by 48 kg per year per 1 ha). In addition to these bacteria living freely in the soil, the nodule growths of some plants of the leguminous family (Leguminosae) were found to contain bacteria (Bacillus radicicola) symbiotically associated with them, which are also able to absorb free atmospheric nitrogen and transfer this nitrogen bound by them to their “host plant”. As you know, this property of leguminous plants (lupin, vetch, seradella, etc.) is widely used to enrich the soil with nitrogenous substances, being a kind of soil fertilization method for subsequent crops of cereals in a plot with plowed and decomposed in the soil, previously grown on it, fertilizing plants. However, these natural sources of replenishment of bound nitrogen in nature can in no way make up for its loss, especially in view of the enormous waste of bound nitrogen in all processes of the destruction of nitrogenous compounds in fuel, as well as when nitrogenous explosives are used. Taking into account the need for nitrogenous food of the world's population, estimated at 1.6 billion people, and the annual growth of the world's population in countries with statistics alone, of 4 million people. or 400 million per century, this loss of bound nitrogen in nature must be considered very significant. William Crooks sounded the alarm back in 1898, predicting the death of mankind from starvation in the near future, when, according to his calculations, the only rich deposits of Chilean saltpeter on the globe, the resource of bound nitrogen, which Ch. arr. was supposed to fill the urgent need of agriculture in nitrogen fertilizers, but instead was rapaciously squandered for military purposes, since most explosives were made by the action of nitric acid obtained from Chilean saltpeter. Indeed, although Crookes somewhat underestimated the reserves of saltpeter in Chile, however, according to the latest geological calculations, even if we accept only the pre-war norm for the production of Chilean saltpeter (2,750,000 tons of saltpeter with a content of 400,000 tons of bound nitrogen), its reserves (600 million tons) tons of saltpeter containing 30 million tons of bound nitrogen) cannot last more than 150-200 years (see Saltpeter). However, the reserves of Chilean saltpeter are by no means the only source from which humanity draws its replenishment of the bound nitrogen necessary for its nutrition and industry. According to the data of the International Agricultural Institute in Rome, calculated on the basis of information about the harvests of all countries of the world, the world consumption of fixed nitrogen in 1924 is determined by the amount of about 7,000,000 tons of bound nitrogen; of these, man was able to work out and return to nature only about 1/6 of the part, that is, about 1,200,000 tons of bound nitrogen. In 1924, only 420,000 tons of Chilean nitrate accounted for this amount. The rest of the amount of bound nitrogen entered the general economy of nature to a large extent due to the same natural resources of bound nitrogen in nature as saltpeter, requiring, however, from the side of man some processing. Such natural resources of bound nitrogen include the world's reserves of coal and peat. Hard coal contains, even in poor grades, from 0.5 to 2% of bound nitrogen. The same varieties that are used for the production of coke and lighting gas usually contain from 1.2 to 1.9%, on average 1.3% of bound nitrogen. According to modern geological data, the world reserves of coal should be estimated at an approximate figure of about 8000 billion tons. Considering the content of bound nitrogen in coal at 1%, we get the content of bound nitrogen in the world reserve of coal at 80 billion tons, i.e., in 2000 times more than the content of bound nitrogen in stocks of Chilean saltpeter. This amount could provide mankind's need for fixed nitrogen for 6,000 years if, using coal, it was possible to utilize all the bound nitrogen contained in it. The pre-war annual production of hard coal was 1,350 million tons with a bound nitrogen content (1.3%) of 17 million tons (corresponding to 85 million tons of ammonium nitrate, worth more than 25 billion francs). However, almost all of this amount of bound nitrogen was released into the air as free nitrogen during the combustion of coal in the furnaces of factories, steam locomotives, in home furnaces, etc. Only about 1/50 of this amount was captured by the nitrogen industry and served to produce sulfuric acid ammonium, which is still the most significant, along with saltpeter, a resource for artificial nitrogen fertilizers (Matignon). On average, 12 kg of ammonium sulphate per t. The utilization of fixed nitrogen from peat is not yet a major factor in the economy of fixed nitrogen. That. the use of coal nitrogen only partially alleviates the acute shortage of bound nitrogen for agriculture and industry, but by no means is a solution to the nitrogen problem as a whole. The final solution of this problem was brought with them by science and technology, ch. arr. during the current century, having carried out the fixation of atmospheric nitrogen by technical means. This fixation is carried out mainly by three main methods: 1) by burning nitrogen in the air under the action of a voltaic arc, with the production of nitrogen oxides and nitric acid; this method, due to the endothermic reaction of the N 2 + O 2 compound, requires the expenditure of significant amounts of heat, high voltage, and is cost-effective only if cheap hydroelectric energy is available; 2) by adding nitrogen at high temperature electric oven to calcium carbide, with the formation of calcium cyanamide; the latter either directly goes for fertilizer purposes, or, under the action of water, forms ammonia, which is neutralized to ammonium sulfate or nitrate; 3) by direct connection of atmospheric nitrogen with hydrogen, with the formation of synthetic ammonia; this method (Haber-Bosch) is undoubtedly the greatest achievement of chemical technology in the past part of the 20th century. and one of the greatest achievements of science and technology in the history of mankind.

Despite the fact that in order to increase the yield, it is also necessary to introduce other fertilizers into the soil - phosphorus and potash, yet it is precisely nitrogen fertilizers that play a predominant role in the agricultural economy. If, for example, in meat phosphoric anhydride and potassium oxide contains 0.4% each, then the amount of bound nitrogen in the same product reaches about 3%, i.e. for 30 hours of bound nitrogen in meat there are only 4 hours P 2 O 6 and K 2 O. At the same time, the prices of these three types of artificial fertilizers in 1913, under normal, comparatively, pre-war conditions, were expressed in the following figures: for 1 kg of bound nitrogen - 1.5 francs, and for 1 kg of K 2 O or P 2 O 5 - 0.4 fr. for every. That. we can consider that nitrogen fertilizers give an economic effect 32 times more significant than the effect of the other two classes of fertilizer fertilizers. How significant the role of nitrogen fertilizers is can be seen from the fact that the introduction of artificial nitrogen fertilizers into the soil causes, ceteris paribus, an increase in yield per 1 ton of bound nitrogen applied: for cereals - 20 tons, for potatoes - 200 tons and for beets - 300 tons. To quantify the role of nitrogenous fertilizers introduced into the agricultural economy, it is interesting to at least approximately calculate the total world capital of bound nitrogen involved in the organic life of our planet. With a land surface of the globe of 135,000,000 km 2 and a layer of arable land of 0.4 m, we can estimate (taking the density of the soil as a unit) the entire capital of the entire fertile soil of the earth at 54 billion tons. The average content of bound nitrogen in the soil does not exceed 0.1%. Reducing the whole calculation to 3/4 due to the inclusion of deserts, glaciers, rocks and other barren soils that do not contain nitrogen, we can estimate the total tonnage of bound nitrogen in the soil of the entire globe at about 40 billion tons, i.e., half of all reserves bound nitrogen present in coal, the utilization of which is possible only to the most limited extent.

The world agriculture demand for nitrogen fertilizers is characterized by the following figures (Partington, The Nitrogen Industry):

World consumption of Chilean saltpeter during the war years is not very indicative, because it was affected by the factors of the blockade, difficult transport, etc.

The world production of fixed nitrogen reached 1,200,000 tons per year, of which: about 30% - 360,000 tons were emitted during coking and gasification from hard coal, about 35% - 420,000 tons were produced in the form of Chilean nitrate, about 35% - 420,000 tons were produced by fixing atmospheric nitrogen. In the most recent years, this ratio has somewhat changed in terms of an increase in the production of saltpeter (up to 36.5%) due to a decrease in the utilization of coal nitrogen (about 30%).

Of all the production of bound nitrogen by fixing atmospheric nitrogen, in turn, 60% d. b. attributed to synthetic ammonia, 30% to cyanamide and only 10% to Norwegian synthetic nitrate. A particularly rapid development of the nitrogen industry is observed in Germany, which is characterized by the following figures: in total, nitrogen products were produced in Germany: in 1915 - 64,000 tons of bound nitrogen, in 1919 - 132,000 tons, in 1920 - 190,000 tons, in 1922 g. - 238,000 tons (these quantities do not include imported Chilean saltpeter). The following diagram graphically depicts the extent to which, in 1925, the world demand for fixed nitrogen was met by the mining and processing nitrogen industry.

Of the total amount of bound nitrogen produced, 83% (about 1,000,000 tons) was used for fertilizer, as a result of which an increase in agricultural products was obtained, equivalent to 20,000,000 tons (1.2 billion poods) of wheat, i.e., almost twice more than the entire annual grain export of Russia in the pre-war years. The development of the synthetic nitrogen industry is illustrated by the following figures:

For individual countries, the world productive capacity of plants producing fixed nitrogen compounds in 1925 is subdivided as follows (in tons):

That. in the technical fixation of atmospheric nitrogen by one method or another, Germany is 60%, France - 14%, England - 2.5%, Italy - 4.3%, Japan - 1.9% and the USA - 18%. But the synthetic nitrogen industry is developing extremely rapidly. Already at the present time part of the construction is being completed, and partly a number of new installations are in operation. When all of them begin to function, the total production of synthetic bound nitrogen will be even greater.

Of all the synthetic methods of atmospheric nitrogen fixation, the predominant importance and the greatest prospects should be recognized for the methods for obtaining synthetic ammonia. The main advantage of this way of fixing atmospheric nitrogen is the very insignificant energy consumption for its production, because the energy, in view of the exothermicity of the process, should be. spent, with the rational use of the heat of the reaction itself, exclusively for the compression of gases to a pressure of 200 atm or more. Parsons (Journal of Ind. a. Eng. Chem., v. 9, p. 839, 1917) gives an interesting calculation of the energy expended per ton of bound nitrogen by various methods:

The current state of the synthetic ammonia industry (as of 1925) is characterized by the following figures:

That. 93% of all synthetic ammonia is produced in Germany. When all atmospheric nitrogen fixation plants are completed, the amount of synthetic ammonia produced will be approximately equal, in terms of a ton of bound nitrogen:

In general, all types of technical fixation of atmospheric nitrogen (ammonia, arc process and cyanamide method) will be able to give an annual production, probably somewhat less than the above, namely:

About 7,400 tons of concentrated ammonia water containing about 400 tons of bound nitrogen was produced in the USSR in 1924; in addition, a significant amount of Chilean nitrate containing 1,700 tons of bound nitrogen was imported. One can get an idea of ​​the needs of the USSR from the following figures. During the war, Russia spent about 330,000 tons of saltpeter with 48,000 tons of bound nitrogen on the production of explosives. The need for nitrogenous fertilizers for crops of sugar beet, cotton and other industrial plants amounts to tens of thousands of tons, and the need for fertilizers for peasant farms - many hundreds of thousands of tons of bound nitrogen. The lack of fertilizers causes a weak harvest in the USSR, on average, 6.5 centners of bread and 98 centners of sugar beet per 1 hectare, against 24.5 centners of bread and 327.5 centners of sugar beet in Western European countries that use nitrogen and other artificial fertilizers (Moser). Resolute measures are now being taken in the USSR to ensure the development of the nitrogen industry. Cm. .

Nitrogen (N 2) was discovered by J. Priestley in 1774. The name "nitrogen" in Greek means "lifeless". It is due to the fact that nitrogen does not support the processes of combustion and respiration. But for all the basic life processes of plant and living organisms, nitrogen is extremely important.

Element characteristic

7 N 1s 2 2s 2 2p 3

Isotopes: 14 N (99.635%); 15 N (0.365%)

Clark in the earth's crust 0.01% by weight. In the atmosphere, 78.09% by volume (75.6% by mass). Nitrogen is a part of living matter (proteins, nucleic acids, and other organic matter). In the hydrosphere, nitrogen is present in the form of nitrates (NO 3). Nitrogen atoms are the 5th most abundant in the universe.

The most important N-containing inorganic substances.

Free (molecular) nitrogen

Nitrogen atoms are interconnected by three covalent non-polar bonds: one of them is a sigma bond, 2 are pi bonds. The breaking energy is very high.

Physical properties

At normal temperature and atmospheric pressure, N 2 is a colorless gas, odorless and tasteless, slightly lighter than air, very poorly soluble in water. It is transferred to a liquid state with great difficulty (Tbp -196 "C). Liquid nitrogen has a high heat of vaporization and is used to create low temperatures (refrigerant).

How to get

Nitrogen is present in the air in a free state, so the industrial method of obtaining is to separate the air mixture (rectification of liquid air).

Under laboratory conditions, small amounts of nitrogen can be obtained in the following ways:

1. Passing air over hot copper, which absorbs oxygen due to the reaction: 2Cu + O 2 \u003d 2CiO. What remains is nitrogen with impurities of inert gases.

2. Redox decomposition of some ammonium salts:

NH 4 NO 2 \u003d N 2 + 2H 2 O

(NH 4) 2 Cr 2 O 7 \u003d N 2 + Cr 2 O 3 + 4H 2 O

3. Oxidation of ammonia and ammonium salts:

4NH 3 + 3O 2 \u003d 2N 2 + 6H 2 O

8NH 3 + ZBr 2 = N 2 + 6NH 4 Br

NH 4 Cl + NaNO 2 \u003d N 2 + NaCl + 2H 2 O

Chemical properties

Molecular nitrogen is a chemically inert substance due to the exceptionally high stability of N 2 molecules. Only the reactions of combination with metals proceed more or less easily. In all other cases, to initiate and accelerate reactions, it is necessary to use high temperatures, spark electric discharges, ionizing radiation, catalysts (Fe, Cr, V, Ti and their compounds).

Reactions with reducing agents (N 2 - oxidizing agent)

1. Interaction with metals:

The reactions of formation of alkali and alkaline earth nitrides Me proceed both with pure nitrogen and during the combustion of metals in air

N 2 + 6Li = 2Li 3 N

N 2 + 6Cs = 2Cs 3 N

N 2 + 3Mg \u003d Mg 3 N 2

2. Interaction with hydrogen (the reaction is of great practical importance):

N 2 + ZN 2 \u003d 2NH 3 ammonia

3. Interaction with silicon and carbon

2N 2 + 3Si \u003d Si 3 N 4 silicon (IV) nitride

N 2 + 2C \u003d (CN) 2 dicyano

2N 2 + 5C + 2Na 2 CO 3 \u003d 4NaCN + 3CO 2 sodium cyanide

Reactions with oxidizing agents (N 2 - reducing agent)

These reactions do not proceed under normal conditions. Nitrogen does not directly interact with fluorine and other halogens, but the reaction with oxygen occurs at the temperature of electric spark discharges:

N 2 + O 2 \u003d 2NO

The reaction is highly reversible; the straight line flows with the absorption of heat (endothermic).

Nitrogen compounds - saltpeter, nitric acid, ammonia - were known long before nitrogen was obtained in a free state. In 1772, D. Rutherford, burning phosphorus and other substances in a glass bell, showed that the gas remaining after combustion, which he called "suffocating air", does not support breathing and combustion. In 1787, A. Lavoisier established that the "vital" and "suffocating" gases that make up the air are simple substances, and proposed the name "Nitrogen". In 1784, G. Cavendish showed that nitrogen is part of saltpeter; this is where the Latin name Azot comes from (from the late Latin nitrum - saltpeter and the Greek gennao - I give birth, I produce), proposed in 1790 by J. A. Chaptal. By the beginning of the 19th century, the chemical inertness of nitrogen in the free state and its exceptional role in compounds with other elements as bound nitrogen were clarified. Since then, the "binding" of nitrogen in the air has become one of the most important technical problems in chemistry.

Distribution of nitrogen in nature. Nitrogen is one of the most common elements on Earth, and most of it (about 4 10 15 tons) is concentrated in the free state in the atmosphere. In the air, free nitrogen (in the form of N 2 molecules) is 78.09% by volume (or 75.6% by mass), not counting minor impurities in the form of ammonia and oxides. The average nitrogen content in the lithosphere is 1.9·10 -3% by weight. Natural Nitrogen compounds are ammonium chloride NH 4 Cl and various nitrates. Large accumulations of saltpeter are characteristic of a dry desert climate (Chile, Central Asia). For a long time saltpeter was the main supplier of nitrogen for industry (now the industrial synthesis of ammonia from atmospheric nitrogen and hydrogen is of primary importance for the binding of nitrogen). Small amounts of bound nitrogen are found in coal (1-2.5%) and oil (0.02-1.5%), as well as in the waters of rivers, seas and oceans. Nitrogen accumulates in soils (0.1%) and in living organisms (0.3%).

Although the name "Nitrogen" means "non-life-sustaining", it is in fact an essential element for life. The protein of animals and humans contains 16-17% nitrogen. In the organisms of carnivorous animals, protein is formed due to the consumed protein substances that are present in the organisms of herbivorous animals and in plants. Plants synthesize protein by assimilating nitrogenous substances contained in the soil, mainly inorganic ones. This means that amounts of nitrogen enter the soil due to nitrogen-fixing microorganisms capable of converting free nitrogen from the air into nitrogen compounds.

The nitrogen cycle occurs in nature leading role in which microorganisms play - nitrifying, denitrifying, nitrogen-fixing and others. However, as a result of the extraction of a huge amount of bound nitrogen from the soil by plants (especially in intensive agriculture), the soils turn out to be depleted in nitrogen. Nitrogen deficiency is typical for agriculture in almost all countries, nitrogen deficiency is also observed in animal husbandry ("protein starvation"). On soils poor in available nitrogen, plants develop poorly. nitrogen fertilizers and protein feeding of animals - essential tool the rise of agriculture. Human economic activity disrupts the cycle of nitrogen. Thus, fuel combustion enriches the atmosphere with nitrogen, and plants that produce fertilizers bind nitrogen in the air. Transportation of fertilizers and agricultural products redistributes nitrogen on the earth's surface. Nitrogen is the fourth most abundant element in the solar system (after hydrogen, helium and oxygen).

Isotopes, atom and molecule of nitrogen. Natural nitrogen consists of two stable isotopes: 14 N (99.635%) and 15 N (0.365%). The 15 N isotope is used in chemical and biochemical research as a labeled atom. Of the artificial radioactive isotopes of nitrogen, 13 N has the longest half-life (T ½ = 10.08 min), the rest are very short-lived. In the upper atmosphere, under the action of neutrons from cosmic radiation, 14 N is converted into a radioactive isotope of carbon 14 C. This process is also used in nuclear reactions to obtain 14 C. The outer electron shell of the nitrogen atom consists of 5 electrons (one lone pair and three unpaired ones - configuration 2s 2 2p 3. Most often, nitrogen in compounds is 3-covalent due to unpaired electrons (as in ammonia NH 3). pairs of electrons can lead to the formation of another covalent bond, and Nitrogen becomes 4-covalent (as in the ammonium ion NH 4).The oxidation states of Nitrogen change from +5 (in N 2 O 5) to -3 (in NH 3). Under normal conditions, in the free state, nitrogen forms an N 2 molecule, where the N atoms are linked by three covalent bonds. The nitrogen molecule is very stable: its dissociation energy into atoms is 942.9 kJ / mol (225.2 kcal / mol), therefore, even at t approx. 3300°C, the degree of dissociation of nitrogen is only about 0.1%.

Physical properties of nitrogen. Nitrogen is slightly lighter than air; density 1.2506 kg / m 3 (at 0 ° C and 101325 n / m 2 or 760 mm Hg), t pl -209.86 ° C, t bp -195.8 ° C. Nitrogen liquefies with difficulty: its critical temperature is rather low (-147.1°C) and its critical pressure is high, 3.39 MN/m 2 (34.6 kgf/cm 2); the density of liquid nitrogen is 808 kg/m 3 . Nitrogen is less soluble in water than oxygen: at 0°C, 23.3 g of nitrogen dissolves in 1 m 3 H 2 O. Better than water, nitrogen is soluble in some hydrocarbons.

Chemical properties of nitrogen. Only with such active metals as lithium, calcium, magnesium, nitrogen interacts when heated to relatively low temperatures. Nitrogen reacts with most other elements at high temperatures and in the presence of catalysts. Nitrogen compounds with oxygen N 2 O, NO, N 2 O 3 , NO 2 and N 2 O 5 are well studied. Of these, upon direct interaction of the elements (4000°C), oxide NO is formed, which, upon cooling, is easily oxidized further to oxide (IV) NO 2 . In air, nitrogen oxides are formed during atmospheric discharges. They can also be obtained by the action of ionizing radiation on a mixture of nitrogen and oxygen. When nitrous N 2 O 3 and nitric N 2 O 5 anhydrides are dissolved in water, nitrous acid HNO 2 and nitric acid HNO 3 are obtained, respectively, forming salts - nitrites and nitrates. Nitrogen combines with hydrogen only at high temperature and in the presence of catalysts, and ammonia NH 3 is formed. In addition to ammonia, numerous other nitrogen-hydrogen compounds are also known, for example, hydrazine H 2 N-NH 2, diimide HN=NH, nitric acid HN 3 (H-N=N≡N), octazone N 8 H 14 and others; most nitrogen compounds with hydrogen have been isolated only in the form of organic derivatives. Nitrogen does not directly interact with halogens, therefore all nitrogen halides are obtained only indirectly, for example, nitrogen fluoride NF 3 - by reacting fluorine with ammonia. As a rule, nitrogen halides are low-resistant compounds (with the exception of NF 3); Nitrogen oxyhalides - NOF, NOCl, NOBr, NO 2 F and NO 2 Cl are more stable. Nitrogen does not combine directly with sulfur either; nitrogenous sulfur N 4 S 4 is obtained by the reaction of liquid sulfur with ammonia. When hot coke reacts with nitrogen, cyanogen (CN) 2 is formed. By heating nitrogen with acetylene C 2 H 2 to 1500°C, hydrogen cyanide HCN can be obtained. The interaction of nitrogen with metals at high temperatures leads to the formation of nitrides (for example, Mg 3 N 2).

When ordinary Nitrogen is exposed to electrical discharges [pressure 130-270 N / m 2 (1-2 mm Hg)] or during the decomposition of B, Ti, Mg and Ca nitrides, as well as during electrical discharges in air, active Nitrogen can be formed , which is a mixture of nitrogen molecules and atoms with an increased energy reserve. Unlike molecular nitrogen, active nitrogen interacts very vigorously with oxygen, hydrogen, sulfur vapor, phosphorus, and certain metals.

Nitrogen is a part of very many important organic compounds (amines, amino acids, nitro compounds, and others).

Getting Nitrogen. In the laboratory, nitrogen can be easily obtained by heating a concentrated solution of ammonium nitrite: NH 4 NO 2 = N 2 + 2H 2 O. The technical method for obtaining nitrogen is based on the separation of preliminarily liquefied air, which is then distilled.

The use of nitrogen. The main part of the extracted free nitrogen is used for the industrial production of ammonia, which is then processed in significant quantities into nitric acid, fertilizers, explosives, etc. In addition to the direct synthesis of ammonia from elements, the cyanamide method developed in 1905 is of industrial importance for the binding of air nitrogen. , based on the fact that at 1000 ° C calcium carbide (obtained by heating a mixture of lime and coal in an electric furnace) reacts with free nitrogen: CaC 2 + N 2 \u003d CaCN 2 + C. The resulting calcium cyanamide decomposes with the release of superheated water vapor ammonia: CaCN 2 + 3H 2 O \u003d CaCO 3 + 2NH 3.

Free nitrogen is used in many industries: as an inert medium in a variety of chemical and metallurgical processes, for filling free space in mercury thermometers, for pumping flammable liquids, etc. Liquid nitrogen is used in various refrigeration plants. It is stored and transported in steel Dewar vessels, gaseous nitrogen in compressed form - in cylinders. Many nitrogen compounds are widely used. The production of bound nitrogen began to develop intensively after World War I and has now reached enormous proportions.

nitrogen in the body. Nitrogen is one of the main biogenic elements that make up essential substances living cells - proteins and nucleic acids. However, the amount of Nitrogen in the body is small (1-3% by dry weight). Molecular nitrogen in the atmosphere can only be assimilated by certain microorganisms and blue-green algae.

Significant reserves of nitrogen are concentrated in the soil in the form of various mineral (ammonium salts, nitrates) and organic compounds (nitrogen of proteins, nucleic acids and their decay products, that is, not yet completely decomposed remains of plants and animals). Plants absorb nitrogen from the soil both in the form of inorganic and some organic compounds. V natural conditions For plant nutrition, soil microorganisms (ammonifiers) are of great importance, which mineralize soil organic nitrogen to ammonium salts. Nitrate nitrogen in the soil is formed as a result of the activity of nitrifying bacteria discovered by S. N. Vinogradsky in 1890, which oxidize ammonia and ammonium salts to nitrates. Part of the nitrate nitrogen assimilated by microorganisms and plants is lost, turning into molecular nitrogen under the action of denitrifying bacteria. Plants and microorganisms assimilate well both ammonium and nitrate nitrogen, reducing the latter to ammonia and ammonium salts. Microorganisms and plants actively convert inorganic ammonium nitrogen into organic nitrogen compounds - amides (asparagine and glutamine) and amino acids. As shown by D. N. Pryanishnikov and V. S. Butkevich, nitrogen is stored and transported in plants in the form of asparagine and glutamine. When these amides are formed, ammonia is neutralized, high concentrations of which are toxic not only for animals, but also for plants. Amides are part of many proteins in both microorganisms and plants, as well as in animals. Synthesis of glutamine and asparagine by enzymatic amidation of glutamvic and aspartic acids is carried out not only in microorganisms and plants, but also in animals within certain limits.

The synthesis of amino acids occurs by reductive amination of a number of aldehyde and keto acids resulting from the oxidation of carbohydrates, or by enzymatic transamination. The end products of the assimilation of ammonia by microorganisms and plants are proteins that are part of the protoplasm and nucleus of cells, as well as deposited in the form of storage proteins. Animals and humans are only able to synthesize amino acids to a limited extent. They cannot synthesize eight essential amino acids (valine, isoleucine, leucine, phenylalanine, tryptophan, methionine, threonine, lysine), and therefore for them the main source of nitrogen is proteins consumed with food, that is, ultimately, plant proteins and microorganisms.

Proteins in all organisms undergo enzymatic breakdown, the end products of which are amino acids. At the next stage, as a result of deamination, the organic nitrogen of amino acids is again converted into inorganic ammonium nitrogen. In microorganisms, and especially in plants, ammonium nitrogen can be used for new synthesis of amides and amino acids. In animals, the neutralization of ammonia formed during the breakdown of proteins and nucleic acids is carried out by the synthesis of uric acid (in reptiles and birds) or urea (in mammals, including humans), which are then excreted from the body. From the point of view of nitrogen metabolism, plants, on the one hand, and animals (and humans), on the other, differ in that in animals, the utilization of the resulting ammonia is carried out only to a weak extent - most of it is excreted from the body; in plants, nitrogen exchange is "closed" - the nitrogen that enters the plant returns to the soil only together with the plant itself.