Nitrogen in nature. Nitrogen - Great Soviet Encyclopedia Physical properties of nitrogen 2

Nitrogen

NITROGEN-A; m.[French azote from Greek. an- - not-, without- and zōtikos - giving life]. Chemical element (N), a colorless and odorless gas that does not support respiration and combustion (it makes up the bulk of the air by volume and mass, and is one of the main elements of plant nutrition).

◁ Nitrogen, oh, oh. A-th acid. A fertilizers. Nitrogenous, oh, oh. A-th acid.

nitrogen

(lat. Nitrogenium), chemical element of group V of the periodic table. Name from Greek. a... is a negative prefix, and zōē is life (does not support breathing and combustion). Free nitrogen consists of 2-atomic molecules (N 2); colorless and odorless gas; density 1.25 g/l, t pl –210ºC, t kip –195.8ºC. Chemically very inert, but reacts with complex compounds of transition metals. The main component of air (78.09% of the volume), the separation of which produces industrial nitrogen (more than 3/4 goes to the synthesis of ammonia). Used as an inert medium for many technological processes; liquid nitrogen is a refrigerant. Nitrogen is one of the main biogenic elements that is part of proteins and nucleic acids.

NITROGEN

NITROGEN (lat. Nitrogenium - giving rise to nitrate), N (read “en”), a chemical element of the second period of the VA group of the periodic system, atomic number 7, atomic mass 14.0067. In its free form, it is a colorless, odorless, and tasteless gas; it is poorly soluble in water. Consists of diatomic N 2 molecules with high strength. Refers to non-metals.
Natural nitrogen consists of stable nuclides (cm. NUCLIDE) 14 N (content in the mixture 99.635% by weight) and 15 N. Configuration of the outer electronic layer 2 s 2 2p 3 . The radius of the neutral nitrogen atom is 0.074 nm, the radius of the ions: N 3- - 0.132, N 3+ - 0.030 and N 5+ - 0.027 nm. The sequential ionization energies of the neutral nitrogen atom are, respectively, 14.53, 29.60, 47.45, 77.47 and 97.89 eV. According to the Pauling scale, the electronegativity of nitrogen is 3.05.
History of discovery
Discovered in 1772 by the Scottish scientist D. Rutherford in the composition of the combustion products of coal, sulfur and phosphorus as a gas unsuitable for breathing and combustion (“suffocating air”) and, unlike CO 2, not absorbed by an alkali solution. Soon the French chemist A.L. Lavoisier (cm. LAVOISIER Antoine Laurent) came to the conclusion that the “suffocating” gas is part of the atmospheric air, and proposed the name “azote” for it (from the Greek azoos - lifeless). In 1784, the English physicist and chemist G. Cavendish (cm. CAVENDISH Henry) established the presence of nitrogen in nitrate (hence the Latin name for nitrogen, proposed in 1790 by the French chemist J. Chantal).
Being in nature
In nature, free (molecular) nitrogen is part of the atmospheric air (in air 78.09% by volume and 75.6% by mass of nitrogen), and in bound form - in the composition of two nitrates: sodium NaNO 3 (found in Chile, hence name Chilean saltpeter (cm. CHILEAN SALTPETER)) and potassium KNO 3 (found in India, hence the name Indian saltpeter) - and a number of other compounds. Nitrogen ranks 17th in abundance in the earth's crust, accounting for 0.0019% of the earth's crust by mass. Despite its name, nitrogen is present in all living organisms (1-3% by dry weight), being the most important biogenic element (cm. BIOGENIC ELEMENTS). It is part of the molecules of proteins, nucleic acids, coenzymes, hemoglobin, chlorophyll and many other biologically active substances. Some so-called nitrogen-fixing microorganisms are able to assimilate molecular nitrogen from the air, converting it into compounds available for use by other organisms (see Nitrogen fixation (cm. NITROGEN FIXATION)). The transformation of nitrogen compounds in living cells is the most important part of metabolism in all organisms.
Receipt
In industry, nitrogen is obtained from the air. To do this, the air is first cooled, liquefied, and the liquid air is subjected to distillation. Nitrogen has a slightly lower boiling point (-195.8°C) than the other component of air, oxygen (-182.9°C), so when liquid air is gently heated, nitrogen evaporates first. Nitrogen gas is supplied to consumers in compressed form (150 atm. or 15 MPa) in black cylinders with a yellow “nitrogen” inscription. Store liquid nitrogen in Dewar flasks (cm. DEWARD VESSEL).
In the laboratory, pure (“chemical”) nitrogen is obtained by adding a saturated solution of ammonium chloride NH 4 Cl to solid sodium nitrite NaNO 2 when heated:
NaNO 2 + NH 4 Cl = NaCl + N 2 + 2H 2 O.
You can also heat solid ammonium nitrite:
NH 4 NO 2 = N 2 + 2H 2 O.
Physical and chemical properties
The density of gaseous nitrogen at 0 °C is 1.25046 g/dm 3, liquid nitrogen (at boiling point) is 0.808 kg/dm 3. Nitrogen gas at normal pressure at a temperature of –195.8 °C turns into a colorless liquid, and at a temperature of –210.0 °C into a white solid. In the solid state, it exists in the form of two polymorphic modifications: below –237.54 °C the form with a cubic lattice is stable, above – with a hexagonal lattice.
The critical temperature of nitrogen is –146.95 °C, the critical pressure is 3.9 MPa, the triple point lies at a temperature of –210.0 °C and a pressure of 125.03 hPa, from which it follows that nitrogen at room temperature is not at any, even very high pressure, cannot be turned into liquid.
The heat of evaporation of liquid nitrogen is 199.3 kJ/kg (at boiling point), the heat of fusion of nitrogen is 25.5 kJ/kg (at a temperature of –210 °C).
The binding energy of atoms in the N 2 molecule is very high and amounts to 941.6 kJ/mol. The distance between the centers of atoms in a molecule is 0.110 nm. This indicates that the bond between the nitrogen atoms is triple. The high strength of the N 2 molecule can be explained within the framework of the molecular orbital method. The energy scheme for filling the molecular orbitals in the N 2 molecule shows that only the bonding s- and p-orbitals are filled with electrons. The nitrogen molecule is non-magnetic (diamagnetic).
Due to the high strength of the N 2 molecule, the decomposition processes of various nitrogen compounds (including the notorious explosive RDX (cm. RDX)) when heated, impacted, etc. lead to the formation of N 2 molecules. Since the volume of the resulting gas is much greater than the volume of the original explosive, an explosion occurs.
Chemically, nitrogen is quite inert and at room temperature reacts only with the metal lithium (cm. LITHIUM) with the formation of solid lithium nitride Li 3 N. In compounds it exhibits different oxidation states (from –3 to +5). Forms ammonia with hydrogen (cm. AMMONIA) NH3. Hydrazine is obtained indirectly (not from simple substances) (cm. HYDRAZINE) N 2 H 4 and hydronitric acid HN 3. Salts of this acid are azides (cm. AZIDS). Lead azide Pb(N 3) 2 decomposes on impact, so it is used as a detonator, for example, in cartridge capsules.
Several nitrogen oxides are known (cm. NITROGEN OXIDES). Nitrogen does not react directly with halogens; NF 3 , NCl 3 , NBr 3 and NI 3 , as well as several oxyhalides (compounds that, in addition to nitrogen, contain both halogen and oxygen atoms, for example, NOF 3 ) are obtained indirectly.
Nitrogen halides are unstable and easily decompose when heated (some during storage) into simple substances. Thus, NI 3 precipitates when aqueous solutions of ammonia and iodine tincture are combined. Even with a slight shock, dry NI 3 explodes:
2NI 3 = N 2 + 3I 2.
Nitrogen does not react with sulfur, carbon, phosphorus, silicon and some other non-metals.
When heated, nitrogen reacts with magnesium and alkaline earth metals, resulting in salt-like nitrides of the general formula M 3 N 2, which decompose with water to form the corresponding hydroxides and ammonia, for example:
Ca 3 N 2 + 6H 2 O = 3Ca(OH) 2 + 2NH 3.
Alkali metal nitrides behave similarly. The interaction of nitrogen with transition metals leads to the formation of solid metal-like nitrides of various compositions. For example, when iron and nitrogen interact, iron nitrides of the composition Fe 2 N and Fe 4 N are formed. When nitrogen is heated with acetylene C 2 H 2, hydrogen cyanide HCN can be obtained.
Of the complex inorganic nitrogen compounds, nitric acid is the most important (cm. NITRIC ACID) HNO 3, its salts nitrates (cm. NITRATES), and also nitrous acid HNO 2 and its salts nitrites (cm. NITRITES).
Application
In industry, nitrogen gas is used mainly to produce ammonia (cm. AMMONIA). As a chemically inert gas, nitrogen is used to provide an inert environment in various chemical and metallurgical processes, when pumping flammable liquids. Liquid nitrogen is widely used as a refrigerant (cm. REFRIGERANT), it is used in medicine, especially in cosmetology. Nitrogen mineral fertilizers are important in maintaining soil fertility (cm. MINERAL FERTILIZERS).

Encyclopedic Dictionary. 2009 .

Synonyms:

See what “nitrogen” is in other dictionaries:

- (N) chemical element, gas, colorless, tasteless and odorless; makes up 4/5 (79%) air; beat weight 0.972; atomic weight 14; condenses into liquid at 140 °C. and pressure 200 atmospheres; constituent of many plant and animal substances. Dictionary… … Dictionary of foreign words of the Russian language

NITROGEN- NITROGEN, chemical. element, symbol N (French AZ), serial number 7, at. V. 14.008; boiling point 195.7°; 1 l A. at 0° and 760 mm pressure. weighs 1.2508 g [lat. Nitrogenium (“generating saltpeter”), German. Stickstoff (“suffocating… … Great Medical Encyclopedia

- (lat. Nitrogenium) N, chemical element of group V of the periodic table, atomic number 7, atomic mass 14.0067. The name is from the Greek a negative prefix and zoe life (does not support respiration or combustion). Free nitrogen consists of 2 atomic... ... Big Encyclopedic Dictionary

nitrogen- a m. azote m. Arab. 1787. Lexis.1. alchemist The first matter of metals is metallic mercury. Sl. 18. Paracelsus set off to the end of the world, offering everyone his Laudanum and his Azoth for a very reasonable price, for the healing of all possible... ... Historical Dictionary of Gallicisms of the Russian Language

- (Nitrogenium), N, chemical element of group V of the periodic system, atomic number 7, atomic mass 14.0067; gas, boiling point 195.80 shs. Nitrogen is the main component of air (78.09% by volume), is part of all living organisms (in the human body... ... Modern encyclopedia

Nitrogen- (Nitrogenium), N, chemical element of group V of the periodic system, atomic number 7, atomic mass 14.0067; gas, boiling point 195.80 °C. Nitrogen is the main component of air (78.09% by volume), is part of all living organisms (in the human body... ... Illustrated Encyclopedic Dictionary

- (chemical sign N, atomic weight 14) one of the chemical elements; colorless gas, odorless, tasteless; very little soluble in water. Its specific gravity is 0.972. Pictet in Geneva and Calhet in Paris succeeded in condensing nitrogen by subjecting it to high pressure... Encyclopedia of Brockhaus and Efron

N (lat. Nitrogenium * a. nitrogen; n. Stickstoff; f. azote, nitrogene; i. nitrogeno), chemical. element of group V is periodic. Mendeleev system, at.sci. 7, at. m. 14.0067. Opened in 1772 researcher D. Rutherford. Under normal conditions A.… … Geological encyclopedia

Male, chem. base, main element of saltpeter; saltpeter, saltpeter, saltpeter; it is also the main, in quantity, component of our air (nitrogen 79 volumes, oxygen 21). Nitrogenous, nitrogenous, nitrogenous, containing nitrogen. Chemists distinguish... Dahl's Explanatory Dictionary

Organogen, nitrogen Dictionary of Russian synonyms. nitrogen noun, number of synonyms: 8 gas (55) non-metal... Dictionary of synonyms

Nitrogen is a gas that extinguishes flames because it does not burn and does not support combustion. It is obtained by fractional distillation of liquid air and stored under pressure in steel cylinders. Nitrogen is used mainly for the production of ammonia and calcium cyanamide, and... ... Official terminology

Books

• Chemistry tests Nitrogen and phosphorus Carbon and silicon Metals Grade 9 To the textbook G E Rudzitis F G Feldman Chemistry Grade 9, Borovskikh T.. This manual fully complies with the federal state educational standard (second generation). The manual includes tests covering the topics of the textbook by G. E. Rudzitis, F. G....

Contents 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 in the atmosphere above each square kilometer of the earth's surface there is so much nitrogen that up to 50 million tons of sodium nitrate or 10 million tons of ammonia (a compound of nitrogen with hydrogen) can be obtained from it, and yet this constitutes 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. Fixed nitrogen is part of both organic and inorganic matter. Plant and animal life contain nitrogen bound to carbon and oxygen in proteins. In addition, nitrogen-containing inorganic compounds such as nitrates (NO 3 –), nitrites (NO 2 –), cyanides (CN –), nitrides (N 3 –) and azides (N 3 –) are known and can be obtained in large quantities ).

Historical information.

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. Without establishing the elemental nature of the gas remaining after combustion, Lavoisier called it azote, which means “lifeless” in ancient Greek. 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 "saltpeter-forming".

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 lightning conditions. S. Arrhenius found that up to 400 million tons of nitrogen are fixed annually in this way. 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 nitrogen; reaching the soil, they turn into nitrites and nitrates. Plants use nitrates to form plant proteins. Animals, feeding on these plants, assimilate the protein substances of the plants and convert them into animal proteins. After the death of animals and plants, they decompose and 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 by other bacteria to nitrates. This is how the nitrogen cycle occurs in nature, or the nitrogen cycle.

Structure of the nucleus and electron shells.

There are two stable isotopes of nitrogen in nature: 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 2p y 1 2p z 1. Consequently, the outer (second) electron shell contains 5 electrons that can participate in the formation of 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 determinations of gas density it has been 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 outer 2 p-electrons 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 temperatures, 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. table 1).

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 during 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. Nitrogen is obtained in huge quantities 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 low cost of the method of liquefying purified air (water vapor, CO 2, dust, and other impurities are 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. The noble gases are released first, then nitrogen, and liquid oxygen remains. Purification is achieved by repeated fractionation processes. This method produces many millions of tons of nitrogen annually, mainly for the synthesis of ammonia, which is the feedstock in the production technology 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.

Nitrogen can be obtained in small quantities in the laboratory in various ways by oxidizing ammonia or ammonium ion, for example:

The process of oxidation of ammonium ion with 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:

The catalytic decomposition of ammonia at high temperatures can also produce nitrogen:

Physical properties.

Some physical properties of nitrogen are given in table. 1.

Table 1. SOME PHYSICAL PROPERTIES OF NITROGEN
Density, g/cm 3	0.808 (liquid)
Melting point, °C	–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/(molCH)	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 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 subsequent 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 of 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 associated with it. Therefore, many compounds of other members of the VA subgroup either have no analogues among nitrogen compounds, or similar nitrogen compounds turn out to be unstable. So, PCl 5 is a stable compound, but NCl 5 does not exist. A nitrogen atom is capable of bonding 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, forming partially ionic nitrides M x N y. In these compounds, nitrogen is negatively charged. In table Table 2 shows the oxidation states and examples of corresponding compounds.

Nitrides.

Compounds of nitrogen 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 bonding in these compounds involves the transfer of electrons from the metal to nitrogen to form the N3– ion. Such nitrides include Li 3 N, Mg 3 N 2, Zn 3 N 2 and Cu 3 N 2. Apart from lithium, other alkali metals do not form IA subgroups of nitrides. Ionic nitrides have high melting points and react with water to form NH 3 and metal hydroxides.

Covalent nitrides.

When nitrogen electrons 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 (such as ammonia and hydrazine) are completely 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 temperatures to form an unusual class of compounds in which the nitrogen atoms are distributed among regularly spaced metal atoms. There is no clear electron displacement in these compounds. Examples of such nitrides are Fe 4 N, W 2 N, Mo 2 N, Mn 3 N 2. These compounds are usually completely inert and have good electrical conductivity.

Hydrogen compounds of nitrogen.

Nitrogen and hydrogen react to form compounds vaguely resembling hydrocarbons. The stability of hydrogen nitrates decreases with increasing number of nitrogen atoms in the chain, in contrast to hydrocarbons, which are stable in long chains. The most important hydrogen nitrides are ammonia NH 3 and hydrazine N 2 H 4. These also include hydronitric 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 USA produced approx. 13 million tons of ammonia annually (in terms of anhydrous ammonia).

Molecule structure.

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 lone electron pair is equivalent to the attached group, resulting in the coordination number of nitrogen being 4 and the nitrogen being located at 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 molecular structure. This is explained by the relatively greater strength of intermolecular bonds in water than in ammonia (such intermolecular bonds are called hydrogen bonds).

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 inorganic substances. 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 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 highly electrically conductive. Unbound alkali metal can be separated from such a solution by evaporation of ammonia or the addition of sodium chloride. Solutions of metals in ammonia are good reducing agents. Autoionization occurs in liquid ammonia

similar to the process occurring in water:

Some chemical properties of both systems are compared in Table. 4.

Liquid ammonia as a solvent has an advantage in some cases where it is not possible 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, and K is soluble, and the reaction proceeds completely. In water, such a reaction is impossible due to the rapid interaction of Ca with water.

Obtaining 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-scale ammonia production is also based on the hydrolysis of nitrides, such as Mg 3 N 2, with water. Calcium cyanamide CaCN 2 when interacting with water also forms ammonia. 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 received for the synthesis of ammonia, differing in the process conditions (temperature, pressure, catalyst). There is a method of industrial production through 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 ENVIRONMENT
Aquatic environment	Ammonia environment
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 –
Amphotericity
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 to form the compound NH 3 N 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 predominantly of NH 3, H 2 O and small concentrations of NH 4 + and OH – ions formed during dissociation

The basic 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 (electrophilic compound) will react with NH 3 to form a coordination compound. For example:

Symbol M n+ represents a transition metal ion (B-subgroup of the periodic table, for example, Cu 2+, Mn 2+, etc.). Any protic (i.e. H-containing) acid reacts with ammonia in an 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; it was first used with petroleum fuel (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. Ammonia gas reacts with metals such as Na and K to form amides:

Ammonia also reacts with hydrides and nitrides 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 temperatures, apparently due to hydrogen generated from the decomposition of ammonia into N 2 and H 2:

Hydrogen atoms in the NH 3 molecule can be replaced by 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. When NH 3 reacts with Cl 2, the chloramines NCl 3, NHCl 2 and NH 2 Cl are formed. When ammonia is exposed to sodium hypochlorite NaOCl (formed from NaOH and Cl 2), the end product is hydrazine:

Hydrazine.

The above reactions are a method for producing hydrazine monohydrate with the composition N 2 H 4 P H 2 O. Anhydrous hydrazine is formed by special distillation of the monohydrate with BaO or other water-removing substances. The properties of hydrazine are slightly similar to 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 (such as methylhydrazine) react with oxygen allows it to be used as a component of liquid rocket fuel. 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 scant 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, dimerization occurs completely: 2NO 2 N 2 O 4. In the liquid state, only 1% NO 2 is undimerized, 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 to form nitric acid: 3NO 2 + H 2 O = 2HNO 3 + NO. NO 2 technology is therefore very important as an intermediate stage in the production of an industrially important product - nitric acid.

Nitric oxide(V)

N2O5( outdated. nitric anhydride) is a white crystalline substance obtained by dehydrating 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 expected structure H–O–N=N–O–H.

Nitrous acid

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

Nitrous acid is also formed when an equimolar mixture of NO and NO 2 (or N 2 O 3) is dissolved in water. Nitrous acid is slightly stronger than acetic acid. The oxidation state of nitrogen 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 influence 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 some 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 of the main chemical industry. It is used in the technologies of many other inorganic and organic substances, such as explosives, fertilizers, polymers and fibers, dyes, pharmaceuticals, etc.

Literature:

Nitrogenist's Directory. M., 1969
Nekrasov B.V. Basics of general chemistry. M., 1973
Nitrogen fixation problems. Inorganic and physical chemistry. M., 1982



Nitrogen is a colorless gas, one of the most common chemical elements on our planet; in the periodic table it is designated by the symbol N from lat. Nitrogenum, meaning lifeless (azoos in Greek). In school, we learn that nitrogen gas makes up 78 percent of the earth's atmosphere. If you put it on one pan of an imaginary scale, then on the other pan you would have to pile 4 x 10 15 tons of weights for balance.

Nitrogen in the form of its compounds plays a colossal role in the life of mankind. Farmers annually apply huge amounts of nitrogen fertilizers to the soil. Nitrogen-containing compounds are in increasing demand in industry - these are dyes, various types of fuel, and polymers. It would seem that the need can be easily satisfied due to the vast ocean of the atmosphere. However, every schoolchild is well aware of the inertness of this substance: the diatomic molecules that make up nitrogen gas do not react with virtually any other substances under normal conditions.

At the same time, the circumstance has long been known that forces chemists to persistently search for new ways. This is the biological fixation of nitrogen by some microorganisms, as well as algae, first established by the Russian scientist S. Vinogradsky back in the 90s of the 19th century. It turns out that chemical inertness does not interfere with the absorption of nitrogen by living organisms? After all, they cannot use high temperatures and pressure. This means that among the enzymes - biological catalysts contained in the body of bacteria - there are those that make it possible to convert nitrogen into proteins at ordinary temperatures and pressures in the presence of water and oxygen.

What was striking was that nitrogen-active systems were not unique. Chemists have worked with many of them before and even used them in industrial processes.

Following this, another discovery was made that broke down the psychological barrier regarding nitrogen. As a result, scientists obtained a unique complex of ruthenium and nitrogen: the gas molecule in it was firmly attached to the metal atom. Such complexes of other molecules with metal compounds were previously known and widely studied. However, no one expected that a molecule of “inert” nitrogen could bind so tightly to a metal ion.

Scientists were unable to determine the conditions for the binding of free nitrogen. However, it was found that free nitrogen is also capable of forming complexes with ruthenium compounds, sometimes in the presence of water and oxygen. Then intensive searches began in different countries of the world, and it turned out that nitrogen binds into complexes with a number of different metals.

Here we could again only wonder why neither nitrogen complexes nor its reactions in solutions had been discovered earlier.

Meanwhile, scientists have moved further. Firstly, it was possible to show that the process can be accelerated by using catalysts to bind large quantities of molecular nitrogen. Secondly, it was discovered that under the influence of compounds of the same transition metals, free nitrogen is capable of reacting with certain organic compounds. Thus, a promising way to obtain valuable chemicals from molecular nitrogen was found.

Now it was necessary to connect together two emerging directions - the chemistry of molecular nitrogen complexes and the study of the reaction of its reduction. After all, it was complex formation (as was previously found for other molecules) that, in principle, should have “activated” inert gas molecules. However, in known complexes it remained inert. Long-term theoretical and experimental work provided an answer to the question of what complexes should be like for the nitrogen in them to be chemically active. Naturally, it is impossible to give a detailed description of the developed theory here. But from it, in particular, it follows that complexes active in relation to further reactions can be observed not at ordinary, but at lower temperatures. Scientists began to isolate from solutions a whole set of complexes in which the nitrogen molecule is activated for further reactions.

Encouraged by their success, the researchers tried to bind nitrogen directly in an aqueous solution, using relatively weak reducing agents, just as bacteria and algae do. In search of the missing data, we had to resort to the help of wildlife.

It was already known that in the enzymatic systems of bacteria, the nitrogen molecule activates molybdenum and this metal cannot be replaced by any other metal except vanadium. Researchers focused their attention on compounds of these particular metals, believing that it was not by chance that nature chose them.

In 1970, they finally got the result that researchers had been striving for for many years. It was possible to discover systems that fix nitrogen in the presence of molybdenum and vanadium compounds in aqueous and aqueous-alcoholic media. The primary reaction endpoint appeared to be almost exclusively hydrazine. Under slightly changed conditions, it was possible to observe the predominant formation of ammonia.

So, there is one less paradox in chemistry. The idea of ​​the inertness of nitrogen has been refuted, new ways have been discovered for converting huge atmospheric “deposits” of this gas into products needed by humans.

Nitrogen(from Greek azoos - lifeless, lat. nitrogenium), n, chemical element of group V of the periodic system of Mendeleev, atomic number 7, atomic mass 14.0067; colorless gas, odorless and tasteless.

Historical information. Ammonium compounds—saltpeter, nitric acid, ammonia—were known long before aluminum 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 respiration and combustion. In 1787, A. Lavoisier established that the “vital” and “asphyxiating” gases that make up the air are simple substances, and proposed the name “A.” In 1784, G. Cavendish showed that A. is part of saltpeter; This is where the Latin name A. 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 exclusive role in compounds with other elements as bound nitrogen were elucidated. Since then, the “binding” of air has become one of the most important technical problems of chemistry.

Prevalence in nature. A. is one of the most common elements on Earth, and its bulk (about 4 × 10 15 T) is concentrated in the free state in atmosphere. In the air, free oxygen (in the form of n2 molecules) is 78.09% by volume (or 75.6% by mass), not counting its minor impurities in the form of ammonia and oxides. The average content of aluminum in the lithosphere is 1.9? 10 -3% by weight. Natural compounds of A. - ammonium chloride nh 4 cl and various nitrates. Large accumulations of saltpeter are characteristic of dry desert climates (Chile, Central Asia). For a long time, nitrate was the main supplier of nitrate for industry (now industrial synthesis is of primary importance for the binding of nitrate ammonia from A. air and hydrogen). Small amounts of bound A. are found in coal (1-2.5%) and oil (0.02-1.5%), as well as in the waters of rivers, seas and oceans. A. accumulates in soils (0.1%) and living organisms (0.3%).

Although the name "A." means “non-life-sustaining”, in fact it is an element necessary for life. The protein of animals and humans contains 16 - 17% A. In the organisms of carnivores, protein is formed due to the consumed protein substances present in the organisms of herbivores and plants. Plants synthesize protein by assimilating nitrogenous substances contained in the soil, mainly inorganic. Significant amounts of A. enter the soil thanks to nitrogen-fixing microorganisms capable of converting free A. air into A compounds.

In nature, there is a cycle of nitrogen, in which the main role is played by microorganisms - nitrophying, denitrifying, nitrogen-fixing, etc. However, as a result of the extraction of huge amounts of bound nitrogen from the soil by plants (especially with intensive farming), the soils become depleted of nitrogen. A deficiency is typical for agriculture in almost all countries, there is a deficiency of protein in animal husbandry (“protein starvation”). On soils poor in available A., plants develop poorly. Nitrogen fertilizers and protein feeding of animals is the most important means of boosting agriculture. Human economic activity disrupts the oxygen cycle. Thus, the combustion of fuel enriches the atmosphere of Australia, and factories producing fertilizers bind air. Transportation of fertilizers and agricultural products redistributes oxygen on the surface of the earth.

A. is the fourth most abundant element in the solar system (after hydrogen, helium and oxygen).

Isotopes, atom, molecule. Natural aluminum consists of two stable isotopes: 14 n (99.635%) and 15 n (0.365%). The 15n isotope is used in chemical and biochemical research as labeled atom. Of the artificial radioactive isotopes, A. has the longest half-life 13 n (t 1/2 = 10.08 min) , the rest are very short-lived. In the upper layers of the atmosphere, under the influence of neutrons from cosmic radiation, 14 n turns into the radioactive carbon isotope 14 c. This process is also used in nuclear reactions to produce 14 c. The outer electron shell of an atom consists of 5 electrons (one lone pair and three unpaired - configuration 2 s 2 2 p 3) . Most often, aluminum in compounds is 3-covalent due to unpaired electrons (as in ammonia nh 3). The presence of a lone pair of electrons can lead to the formation of another covalent bond, and A. becomes 4-covalent (as in the ammonium ion nh 4 +). The oxidation states of A. vary from +5 (in n 2 0 5) to -3 (in nh 3). Under normal conditions, in the free state, A. forms a molecule n 2, where the n atoms are connected by three covalent bonds. The A. molecule is very stable: its dissociation energy into atoms is 942.9 kJ/mol (225,2 kcal/mol) , therefore even with t At about 3300°C, the degree of dissociation of A. is only about 0.1%.

Physical and chemical properties. A. slightly lighter than air; density 1.2506 kg/m 3(at 0°C and 101325 n/m 2 or 760 mmHg Art.) , t pl-209.86°С, t kip-195.8°c. A. liquefies with difficulty: its critical temperature is quite low (-147.1 ° C), and its critical pressure is high 3.39 Mn/m 2 (34,6 kgf/cm 2); density of liquid A. 808 kg(m3. In water, A. is less soluble than oxygen: at 0°C in 1 m 3 H 2 O dissolves 23.3 G A. Better than in water, A. soluble in some hydrocarbons.

A. reacts only with active metals such as lithium, calcium, and magnesium when heated to relatively low temperatures. A. reacts with most other elements at high temperatures and in the presence of catalysts. The compounds of A. with oxygen n 2 o, no, n 2 o 3, no 2 and n 2 o 5 have been well studied. From them, with the direct interaction of elements (4000°c), no oxide is formed, which, upon cooling, is easily oxidized further to no 2 dioxide . In the air, aluminum oxides are formed during atmospheric discharges. They can also be obtained by exposing the mixture of oxygen and oxygen to ionizing radiation. When nitrogenous n 2 O 3 and nitrogen n 2 O 5 anhydrides are dissolved in water, respectively, we obtain nitrous acid hno2 and nitric acid hno 3, forming salts - nitrites And nitrates. A. combines with hydrogen only at high temperatures and in the presence of catalysts, and this forms ammonia nh 3. In addition to ammonia, numerous other compounds of ammonia with hydrogen are known, for example hydrazine h 2 n-nh 2, diimide hn=nh, hydronitric acid hn 3 (h-n=n ? n), octazone n 8 h 14, etc.; Most of the compounds of hydrogen with hydrogen are isolated only in the form of organic derivatives. A. does not directly interact with halogens, therefore all A. halides are obtained only indirectly, for example, nitrogen fluoride nf 3 - through the interaction of fluorine with ammonia. As a rule, A. halides are low-resistant compounds (with the exception of nf 3); A. oxyhalides are more stable - nof, noci, nobr, n0 2 f and no2ci. A. also does not directly combine with sulfur; nitrogenous sulfur n 4 s 4 is obtained as a result of the reaction of liquid sulfur with ammonia. When hot coke interacts with alcohol, it forms cyanogen(cn).;. By heating A. with acetylene from 2 h 2 to 1500 ° c can be obtained hydrogen cyanide hcn. The interaction of aluminum with metals at high temperatures leads to the formation nitrides(for example, mg 3 n 2).

When a normal A. is exposed to electrical discharges [pressure 130 - 270 n/m 2(1- 2 mmHg)] or during the decomposition of nitrides B, ti, mg and Ca, as well as during electrical discharges in the air, active aluminum can be formed, which is a mixture of molecules and atoms of aluminum with an increased energy reserve. Unlike molecular, active oxygen interacts very energetically with oxygen, hydrogen, sulfur vapor, phosphorus, and some metals.

A. is part of many important organic compounds ( amines, amino acids, nitro compounds etc.).

Receipt and application. In the laboratory, A. can easily be obtained by heating a concentrated solution of ammonium nitrite: nh4no2 = n 2 + 2h 2 O. The technical method for obtaining A. is based on the separation of pre-liquefied air, which is then subjected to distillation.

The main part of the extracted free ammonia 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, cyanamide, developed in 1905, is of industrial importance for binding ammonia. method 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 A.: CaC + n -= cacn + C. The resulting calcium cyanamide when exposed to superheated water vapor, it decomposes with the release of ammonia:

cacn+ZN 2 O=CaCO 3 +2nh 3 .

Free aluminum is used in many industries: as an inert medium in various chemical and metallurgical processes, for filling free space in mercury thermometers, when pumping flammable liquids, etc. Liquid aluminum is used in various refrigeration units. It is stored and transported in steel Dewar vessels, gaseous A. in compressed form - in cylinders. Many compounds of A are widely used. The production of bound A began to develop rapidly after the 1st World War and has now reached enormous proportions.

Lit.: Nekrasov B.V., Fundamentals of General Chemistry, vol. 1, M., 1965; Remi G., Course of inorganic chemistry, trans. from German, vol. 1, M., 1963: Chemistry and technology of bound nitrogen, [M.-L.], 1934; KHE, vol. 1, M., 1961.

• Designation - N (Nitrogen);
• Period - II;
• Group - 15 (Va);
• Atomic mass - 14.00674;
• Atomic number - 7;
• Atomic radius = 92 pm;
• Covalent radius = 75 pm;
• Electron distribution - 1s 2 2s 2 2p 3 ;
• melting temperature = -209.86°C;
• boiling point = -195.8°C;
• Electronegativity (according to Pauling/according to Alpred and Rokhov) = 3.04/3.07;
• Oxidation state: +5, +4, +3, +2, +1, 0, -1, -2, -3;
• Density (no.) = 0.808 g/cm 3 (-195.8°C);
• Molar volume = 17.3 cm 3 /mol.

Nitrogen compounds:

• Equations of oxidation-reduction reactions of nitrogen...

It is not possible to unambiguously name the scientist who was the first to discover nitrogen for the simple reason that three people did it almost simultaneously in 1772 - Henry Cavendish, Joseph Priestley and Daniel Rutherford (Carl Scheele can also be added to this list). However, not one of the scientists at one time fully understood his discovery. Many people give the "palm" to the Scotsman Daniel Rutherford, since he was the first to publish a master's thesis in which he described the basic properties of "spoiled air."

The actual name was proposed in 1787 by A. Lavoisier.

Nitrogen is the fourth most abundant chemical element in the solar system (after hydrogen, helium and oxygen). Nitrogen is one of the most abundant elements on Earth:

• the earth's atmosphere contains 3.87·10 18 kg of nitrogen - 75.6% (by mass) or 78.08% (by volume);
• the earth's crust contains nitrogen (0.7-1.5) 10 18 kg;
• the earth's mantle contains 1.3·10 19 kg of nitrogen;
• the hydrosphere contains 2·10 16 kg of nitrogen (7·10 14 kg in the form of compounds).

Nitrogen plays a vital role in the life of organisms - it is present in proteins, amino acids, amines, and nucleic acids.

Natural nitrogen consists of two stable isotopes 14 N - 99.635% and 15 N - 0.365%.

The nitrogen atom contains 7 electrons, which are located in two orbitals (s and p) (see Electronic structure of atoms). The inner orbital contains 2 electrons; on the outer one - 5 (one free electron pair + three unpaired electrons, which can form three covalent bonds; see Covalent bond).

When reacting with other chemical elements, the nitrogen atom can exhibit an oxidation state from +5 to -3 (in addition to three valence electrons, another bond can be formed by the donor-acceptor mechanism due to a free electron pair with an atom having a free orbital).

Nitrogen oxidation states:

• +5 - HNO 3;
• +4 - NO 2;
• +3 - HNO 2;
• +2 - NO;
• +1 - N 2 O;
• -1 - NH 2 OH;
• -2 - N 2 H 4 ;
• -3 (most common) - NH 3.

N 2

The three unpaired p-electrons of the nitrogen atom, lying on its outer energy level, have the shape of equal arms of a figure eight, located perpendicular to each other:

When a nitrogen molecule (N2) is formed, the p-orbital located along the X axis of one atom overlaps with a similar p x -orbital of another atom - at the intersection of the orbitals, an increased electron density is formed with the formation of a covalent bond ( σ bond).

Two other orbitals of one atom, located along the Y and Z axes, overlap their lateral surfaces with their “brothers” of another atom, forming two more covalent bonds ( π bonds).

As a result, 3 covalent bonds are formed in the nitrogen molecule (N 2) (two π bonds + one σ bond), i.e., a very strong triple bond appears (see Multiple bonds).

The nitrogen molecule is very strong (dissociation energy 940 kJ/mol) and has low reactivity.

Properties of molecular nitrogen

Under normal conditions, nitrogen is a low-active substance, which is explained by fairly strong interatomic bonds in its molecule, since they are formed by as many as three pairs of electrons. For this reason, nitrogen usually reacts at high temperatures.

• odorless and colorless gas;
• poorly soluble in water;
• soluble in organic solvents;
• can react with metals and non-metals when heated in the presence of a catalyst (under the influence of ionizing radiation);
• nitrogen reacts as an oxidizing agent (with the exception of oxygen and fluorine):
  • Under normal conditions, nitrogen reacts only with lithium:
    6Li + N 2 = 2Li 3 N;
  • When heated, nitrogen reacts with metals:
    2Al + N 2 = 2AlN;
  • at a temperature of 500°C and at high pressure in the presence of iron, nitrogen reacts with hydrogen:
    N 2 + 3H 2 ↔ 2NH 3 ;
  • at a temperature of 1000°C, nitrogen reacts with oxygen, boron, silicon:
    N 2 + O 2 ↔ 2NO.
• nitrogen interacts as a reducing agent:
  • with oxygen:
    N 2 0 +O 2 0 ↔ 2N +2 O -2 (nitric oxide II)
  • with fluorine:
    N 2 0 +3F 2 0 = 2N+3F 3 -1 (nitrogen fluoride III)

Obtaining and using nitrogen

Nitrogen production:

• industrially, nitrogen is obtained by liquefying air with subsequent separation of nitrogen by evaporation;
• laboratory methods for obtaining nitrogen:
  • decomposition of ammonium nitrite:
    NH 4 NO 2 = N 2 + 2H 2 O;
  • reduction of nitric acid with active metals:
    36HNO 3 + 10Fe = 10Fe(NO 3) 3 + 3N 2 + 18H 2 O;
  • decomposition of metal azides (pure nitrogen):
    2NaN 3 → (t) 2Na + 3N 2 ;
  • Atmospheric nitrogen is produced by reacting air with hot coke:
    O 2 + 4N 2 + 2C → 2CO + 4N 2;
  • by passing ammonia over copper (II) oxide at t=700°C:
    2NH 3 + 3CuO → N 2 + 3H 2 O + 3Cu.

Application of nitrogen:

• creation of inert environments in metallurgy;
• synthesis of ammonia and nitric acid;
• production of explosives;
• to create low temperatures;
• production of mineral fertilizers: potassium nitrate (KNO 3); sodium nitrate (NaNO 3); ammonium nitrate (NH 4 NO 3); lime nitrate (Ca(NO 3) 2).