Presentation - nitrogen, its structure and properties. Nitrogen from the atmosphere Model of the electronic structure of the nitrogen atom

DEFINITION

Nitrogen- the seventh element of the Periodic Table. Refers to non-metals. Located in the second period V of group A subgroup.

The serial number is 7. The nuclear charge is +7. Atomic weight - 14.007 amu. There are two isotopes of nitrogen found in nature: 14 N - 99.635% and 15 N - 0.365% (their percentages are indicated in parentheses).

Electronic structure of the nitrogen atom

The nitrogen atom has two shells, like all elements located in the second period. The group number -V - indicates that the outer electronic level of the nitrogen atom contains 5 valence electrons.

Rice. 1. Schematic structure of the nitrogen atom.

The electronic configuration of the ground state is written as follows:

1s 2 2s 2 2p 3 .

Nitrogen is an element of the p-family. The energy diagram for valence electrons in the unexcited state is as follows:

There is no excited state. Based on the number of unpaired electrons, we can say that nitrogen in compounds exhibits valency III. However, valence is also determined by the group number (V), therefore, nitrogen can exhibit two valences - III and V.

Rice. 2. Spatial representation of the structure of the nitrogen atom.

Examples of problem solving

EXAMPLE 1

Nitrogen (N) is a gas whose content in the atmosphere is about 78%. Nitrogen is part of amino acids and nucleotides. The structure of the nitrogen atom determines the physical and chemical properties of the element.

Structure

Nitrogen is the seventh element of the periodic table, located in the fifth group and second period. The relative atomic mass is 14. Under natural conditions, two isotopes of nitrogen are found - 14 N and 15 N.

Rice. 1. Nitrogen in the periodic table.

Nitrogen consists of a nucleus with a charge of +7 and seven electrons distributed over two energy levels. The presence of an element in the fifth group indicates the number of electrons in the outer level and the highest valence. In an unexcited state, there are three electrons at the outer energy level, so nitrogen can exhibit two valences - III and V.

Recording the electronic structure of the nitrogen atom is 1s 2 2s 2 2p 3 or +7 N) 2) 5.

Physical properties

Nitrogen is a diatomic (N 2) gas, odorless and tasteless, poorly soluble in water. Nitrogen can be in gaseous, liquid and solid states. In liquefied form, nitrogen has a boiling point of -196°C. At -209.86°C, nitrogen becomes solid. Under the influence of different temperatures, the crystal lattice of solid nitrogen can change, creating modifications of the element.

Rice. 2. Liquid and solid nitrogen.

Chemical properties

Nitrogen atoms are connected by a triple bond (N ≡ N), which provides maximum strength. Even when nitrogen is heated to 3000°C, slight decomposition of molecules is observed (up to 0.1% of the amount of gas taken). That is why nitrogen is a chemically inactive element. In compounds when heated, nitrogen easily separates from other elements.

The main chemical properties of nitrogen are given in the table.

Compounds of nitrogen with metals and non-metals are called nitrides.

Nitrogen does not react with acids, water and bases. Direct interaction of nitrogen molecules with sulfur and halogens is impossible. Atomic nitrogen reacts more actively with these substances under normal conditions.

Application

Despite the passivity of nitrogen, the element is widely used in industry. In addition, nitrogen is part of cells; without it, the construction of protein and DNA is impossible.

Rice. 3. Nitrogen in DNA.

Nitrogen is used in the production of:

fertilizers;
explosives;
medicines;
dyes;
plastics;
artificial fibers;
ammonia.

Liquid nitrogen is used for cooling, freezing, and oxidizing rocket engines. Nitric oxide is used as an anesthetic and for the production of aerosols.

What have we learned?

We examined the structure of nitrogen, its physical and chemical properties, and applications. Nitrogen consists of a positively charged nucleus and two electron shells containing seven electrons. Nitrogen is a low-active gas. A nitrogen molecule consists of two atoms of the element connected by a triple bond. Nitrogen can be in three states of aggregation. The element reacts with some metals, non-metals and oxygen. Nitrogen is used in industry, medicine, and agriculture. In addition, nitrogen is part of living organisms.

During the lesson you will gain an understanding of the topic "Nitrogen". Get to know nitrogen as a simple substance, ammonia, nitric acid and nitrates. The chemical and physical properties of these substances, the structure of their molecules, and reactions with other substances will be considered. In addition, methods for obtaining these substances by industrial and laboratory methods and their use in various industries will be listed. Consider the properties and uses of nitrous oxide and aqua regia (a combination of three parts hydrochloric acid and one part nitric acid).

Topic: Basic metals and non-metals

Lesson: Nitrogen

1. Electronic structure of the nitrogen atom

The chemical element nitrogen is located in the second period of group 5, the main subgroup. The electron configuration of the nitrogen atom is 1s22s22p3. There are no vacant orbitals at the valence energy level of the nitrogen atom. Consequently, the electron pair of the 2s sublevel cannot be decoupled. See Fig. 1. Therefore, nitrogen cannot be 5-valent. The maximum valency of nitrogen in compounds is 4. In this case, 3 bonds are formed by the exchange mechanism, and one by the donor-acceptor mechanism. Nitrogen exhibits oxidation states from -3 to +5.

For examples of substances with different oxidation states, see Fig. 2.

2. Nitrogen is a simple substance

Allotropy is not typical for nitrogen. It forms one simple substance, N2. It is a molecular substance with a covalent nonpolar bond. The bond is formed using three shared electron pairs, a triple bond - one sigma and 2 pi bonds. The triple bond is very strong. This causes the low reactivity of molecular nitrogen.

Physical properties

Nitrogen is a colorless and odorless gas, poorly soluble in water, slightly lighter than air. Nitrogen reacts with some substances, but the reaction conditions are very harsh (high temperature and pressure, use of a catalyst). Under normal conditions, nitrogen reacts only with lithium, forming lithium nitride.

6Li + N2 = 2Li3N, by hydrolysis of which ammonia can be obtained.

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 (NO2), cyanides (CN), nitrides (N3) and azides (N3). Historical 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. 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 withmass number 15 (contains 7 protons and 8 neutrons). Their ratio is 99.635:0.365, so the atomic mass of nitrogen is 14.008. Unstable isotopes of nitrogen 12 N, 13 N, 16 N, 17 N are obtained artificially.Schematic electronic structure of the nitrogen atom is: 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. it is possible to form compounds with oxidation states from ( II I) to (V), and they are known. see also ATOMIC STRUCTURE.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 ones 2p- electrons of each atom form a triple bond:N:::N:, forming electron pairs. Measured interatomic distance N N equals 1.095 Å . As is the case with 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). Dissociation energy of a mole (28.016 g or 6.023 Ch 10 23 molecules) molecular nitrogen into atoms ( N 2 2N) equal to 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 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: decomposition of azides when heated, decomposition of ammonia with copper(II) oxide, 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/(mol K)	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 densitiesNitrogen liquid and gaseous states are the same. b The amount of energy required to remove the first outer electron and the next one, per 1 mole of atomic nitrogen.

Table 2. OXIDATION STATES OF NITROGEN AND CORRESPONDING COMPOUNDS
Oxidation state	Connection examples
	Ammonia NH 3, ammonium ion NH 4 +, nitrides M 3 N 2
	Hydrazine N2H4
	Hydroxylamine NH 2 OH
	Sodium hyponitrite Na 2 N 2 O 2 , nitric oxide (I) N 2 O
	Nitrogen(II) oxide NO
	Nitrogen(III) oxide N 2 O 3, sodium nitrite NaNO 2
	Nitric oxide (IV) NO 2, dimer N 2 O 4
	Nitric oxide(V) N 2 O 5 , Nitric acid HNO3 and its salts (nitrates)

Nitrides. Compounds of nitrogen with more electropositive elements, metals and nonmetals nitrides are similar to carbides and hydrides. They can be divided depending on the nature of the MN 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 N ion 3 . 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 are 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 in reaction with NH 3 at high temperatures form an unusual class of compounds in which nitrogen atoms are distributed between regularly spaced metal atoms. There is no clear electron displacement in these compounds. Examples of such nitrides 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 interact to form compounds vaguely reminiscent of hydrocarbons (see also ORGANIC CHEMISTRY). 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. NH 3 molecule has an almost pyramidal structure. The bond angle HNH is 107° , which is close to the tetrahedral angle 109° . The lone electron pair is equivalent to the attached group, resulting in the coordination number of nitrogen being 4 and 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.

Table 3. SOME PHYSICAL PROPERTIES OF AMMONIA AND WATER
Property
Density, g/cm 3	0.65 (–10° C)	1.00 (4.0° C)
Melting point, °C	–77,7	0
Boiling point, °C	–33,35	100
Critical temperature, °C	132	374
Critical pressure, atm	112	218
Enthalpy of vaporization, J/g	1368 (–33° C)	2264 (100° C)
Melting enthalpy, J/g	351 (–77° C)	334 (0° C)
Electrical conductivity	5H 10 –11 (–33° C)	4H 10 –8 (18° C)

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 schemeThe 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.

Production of 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 , water. Calcium cyanamide CaCN 2 When interacting with water, it also forms ammonia. The main industrial method for producing ammonia is its catalytic synthesis from atmospheric nitrogen and hydrogen at high temperature and pressure:

Table 4. COMPARISON OF REACTIONS IN WATER AND AMMONIA ENVIRONMENT
Water 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 NH compound 3 H H 2 O, which is often mistakenly considered ammonium hydroxide NH 4 OH; in reality existence NH4OH not proven in solution. An aqueous solution of ammonia (“ammonia”) consists predominantly of NH 3, H 2 O and low concentrations of NH ions 4 + and OH , formed during dissociationThe basic nature of ammonia is explained by the presence of a nitrogen:NH lone electron pair 3. Therefore NH 3 is a Lewis base that has the highest nucleophilic activity, manifested in the form of association with the proton, or nucleus of the hydrogen atom:Any ion or molecule capable of accepting an electron pair (electrophilic compound) will interact with NH 3 with the formation of a coordination compound. For example:

Symbol M n+ represents a transition metal ion (B subgroups of the periodic table, e.g. Cu 2+ , Mn 2+ andetc.). 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 , ammonium phosphate (NH 4) 3PO 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 (e.g. 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 molecule 3 can be replaced by halogen. Iodine reacts with concentrated NH solution 3 , forming a mixture of substances containing N I 3 . This substance is very unstable and explodes at the slightest mechanical impact. When reacting NH 3 c Cl 2 chloramines NCl 3, NHCl 2 and NH 2 Cl are formed. When ammonia is exposed to sodium hypochlorite NaOCl (formed from NaOH and Cl2 ) the final product is hydrazine:

Hydrazine. The above reactions are a method for preparing hydrazine monohydrate of 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. The properties of hydrazine are slightly similar to hydrogen peroxide H 2 O 2 . Pure Anhydrous Hydrazine 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: N2 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). Nitric oxide (I) N 2 O (dianitrogen monoxide) is obtained from the thermal dissociation of ammonium nitrate:The molecule has a linear structureO is fairly inert at room temperature, but at high temperatures it can support the combustion of easily oxidized materials. N 2 O, known as laughing gas, is used for mild anesthesia in medicine.Nitric oxide(II) NO colorless gas, is one of the products of the catalytic thermal dissociation of ammonia in the presence of oxygen:

NO is also formed during the thermal decomposition of nitric acid or during the reaction of copper with dilute nitric acid:NO can be obtained by synthesis from simple substances (N 2 and O 2 ) at very high temperatures, for example in an electrical discharge. The structure of the NO molecule has one unpaired electron. Compounds with this structure interact with electric and magnetic fields. In the liquid or solid state, the oxide is blue in color because the unpaired electron causes partial association in the liquid state and weak dimerization in the solid state: 2NO N2O2. Nitric oxide (III) N2O3 (nitrogen trioxide) nitrous anhydride: N2O3 + H2O2HNO2. Pure N2O3 can be obtained as a blue liquid at low temperatures (20° C) from an equimolecular mixture of NO and NO 2. N2O3 stable only in the solid state at low temperatures (mp 102.3° C), in liquid and gaseous states it again decomposes into NO and NO 2 . Nitric oxide (IV) NO 2 (nitrogen dioxide) also has an unpaired electron in the molecule ( see above nitric oxide (II)). The structure of the molecule assumes a three-electron bond, and the molecule exhibits the properties of a free radical (one line corresponds to two paired electrons):

obtained by the catalytic oxidation of ammonia in excess oxygen or the oxidation of NO in air:

and also by reactions:

At room temperature NO 2 The gas is dark brown in color and has magnetic properties due to the presence of an unpaired electron. At temperatures below 0°C NO 2 molecule dimerizes into dinitrogen tetroxide, and at 9.3° C dimerization proceeds completely: 2NO2N2O4 . In the liquid state, only 1% NO is undimerized 2, and at 100 ° C remains as a dimer of 10% N 2 O 4 . (or N2O4 ) reacts in warm water to form nitric acid: 3NO 2 + H 2 O = 2HNO 3 + NO. NO 2 technology therefore very important as an intermediate stage in obtaining an industrially important product nitric acid.Nitric oxide (V) N2O5 (outdated. nitric anhydride) white crystalline substance, obtained by dehydration of nitric acid in the presence of phosphorus oxide P 4 O 10:

N2O5 easily dissolves in air moisture, again forming HNO3. Properties of N2O5 determined by equilibrium

N 2 O 5 is a good oxidizing agent, it reacts easily, sometimes violently, with metals and organic compounds and, in a pure state, explodes when heated. Probable structure. When the solution is evaporated, a white explosive is formed with the expected structure HON=NOH. Nitrous acid HNO2 is not exists in its pure form, but 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 is dissolved 2 (or N 2 O 3 ) in water. Nitrous acid is slightly stronger than acetic acid. The oxidation state of nitrogen in it is +3 (its structure HON=O), those. it can be both an oxidizing agent and a reducing agent. Under the influence of reducing agents, it is usually restored 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 are highly soluble in water, except for silver nitrite.

NaNO2 used in the production of dyes.Nitric acid HNO3 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. see also CHEMICAL ELEMENTS.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
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. 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 (NO3-), nitrites (NO2-), cyanides (CN-), nitrides (N3-) and azides (N3-) are known and can be obtained in large quantities.
Historical 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. Without establishing the elemental nature of the gas remaining after combustion, Lavoisier called it azote, which in ancient Greek means “lifeless.” In 1772, D. Rutherford of Edinburgh established that this gas is an element and called it “noxious air.” The Latin name for nitrogen comes from the Greek words nitron and gen, which mean “saltpeter-forming.”
Nitrogen fixation and the nitrogen cycle. The term "nitrogen fixation" refers to the process of fixing atmospheric nitrogen N2. 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 mass number 14 (N contains 7 protons and 7 neutrons) and with mass number 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 12N, 13N, 16N, 17N were obtained artificially. Schematically, the electronic structure of the nitrogen atom is as follows: 1s22s22px12py12pz1. 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.
See also ATOMIC STRUCTURE.
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 NєN (or N2). For two nitrogen atoms, the three outer 2p electrons of each atom form a triple bond:N:::N:, forming electron pairs. The measured N-N interatomic distance is 1.095. As in the case of hydrogen (see 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 a transition temperature a (r) b -237.39 ° C. Modification b melts at -209.96 ° C and boils at -195.78 ° C at 1 atm (see table 1). The dissociation energy of a mole (28.016 g or 6.023 * 10 23 molecules) of molecular nitrogen into atoms (N2 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, CO2, 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 - decomposition of azides when heated, decomposition of ammonia with copper(II) oxide, 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/cm3 0.808 (liquid) Melting point, °C -209.96 Boiling point, °C -195.8 Critical temperature, °C -147.1 Critical pressure, atma 33.5 Critical density, g/cm3 a 0.311 Specific heat, J/(mol) 14.56 (15° C) Pauling electronegativity 3 Covalent radius, 0.74 Crystalline radius, 1.4 (M3-) Ionization potential, Wb

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 electron and the next one, 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 at the 2s level and three half-filled 2p 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 at all, or similar nitrogen compounds turn out to be unstable. So, PCl5 is a stable compound, but NCl5 does not exist. A nitrogen atom is capable of bonding with another nitrogen atom, forming several fairly stable compounds, such as hydrazine N2H4 and metal azides MN3. 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 MxNy. In these compounds, nitrogen is negatively charged. In table Table 2 shows the oxidation states and examples of corresponding compounds.
Table 2. OXIDATION STATES OF NITROGEN AND CORRESPONDING COMPOUNDS
Oxidation state Examples of compounds
-III Ammonia NH3, ammonium ion NH4+, nitrides M3N2 -II Hydrazine N2H4 -I Hydroxylamine NH2OH I Sodium hyponitrite Na2N2O2, nitric oxide(I) N2O II Nitric oxide(II) NO III Nitrogen oxide N2O3, sodium nitrite NaNO2 IV Oxide nitrogen(IV) NO2, dimer N2O4 V Nitrogen oxide (V) N2O5, nitric acid HNO3 and its salts (nitrates) Nitrides. Compounds of nitrogen with more electropositive elements, metals and nonmetals - 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 Li3N, Mg3N2, Zn3N2 and Cu3N2. 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 NH3 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 (NF3 and NCl3). Covalent nitrides include, for example, Si3N4, P3N5 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 NH3 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 Fe4N, W2N, Mo2N, Mn3N2. These compounds are usually completely inert and have good electrical conductivity.
Hydrogen compounds of nitrogen. Nitrogen and hydrogen interact to form compounds that vaguely resemble hydrocarbons (see also ORGANIC CHEMISTRY). 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 NH3 and hydrazine N2H4. These also include hydronitric acid HNNN (HN3).
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 NH3 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 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.

Table 3. SOME PHYSICAL PROPERTIES OF AMMONIA AND WATER

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 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 CaCl2 and K, since CaCl2 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. Production of ammonia. Gaseous NH3 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 Mg3N2, with water. Calcium cyanamide CaCN2 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 through thermal distillation of coal. The names of F. Haber and K. Bosch are associated with the technological development of ammonia synthesis.
Chemical properties of ammonia. In addition to the reactions mentioned in table. 4, ammonia reacts with water to form the compound NH3НH2O, which is often mistakenly considered ammonium hydroxide NH4OH; in fact, the existence of NH4OH in solution has not been proven. An aqueous solution of ammonia (“ammonia”) consists predominantly of NH3, H2O and small concentrations of NH4+ and OH- ions formed during dissociation

The basic nature of ammonia is explained by the presence of a lone electron pair of nitrogen:NH3. Therefore, NH3 is a Lewis base, which has the highest nucleophilic activity, manifested in the form of association with the proton, or nucleus of the hydrogen atom:

Any ion or molecule capable of accepting an electron pair (electrophilic compound) will react with NH3 to form a coordination compound. For example:

The symbol Mn+ represents a transition metal ion (B subgroup of the periodic table, for example, Cu2+, Mn2+, etc.). Any protic (i.e. H-containing) acid reacts with ammonia in an aqueous solution to form ammonium salts, such as ammonium nitrate NH4NO3, ammonium chloride NH4Cl, ammonium sulfate (NH4)2SO4, ammonium phosphate (NH4)3PO4. 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 NH2CONH2, 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, NaNH2) react with N2O when heated, forming azides:

Gaseous NH3 reduces heavy metal oxides to metals at high temperatures, apparently due to hydrogen produced by the decomposition of ammonia into N2 and H2:

Hydrogen atoms in the NH3 molecule can be replaced by halogen. Iodine reacts with a concentrated solution of NH3, forming a mixture of substances containing NI3. This substance is very unstable and explodes at the slightest mechanical impact. The reaction of NH3 with Cl2 produces the chloramines NCl3, NHCl2 and NH2Cl. When ammonia is exposed to sodium hypochlorite NaOCl (formed from NaOH and Cl2), the end product is hydrazine:

Hydrazine. The above reactions represent a method for producing hydrazine monohydrate with the composition N2H4ЧH2O. 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 H2O2. 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: N2O, NO, N2O3, NO2 (N2O4), N2O5. There is scant information on the formation of nitrogen peroxides (NO3, NO4). Nitrogen(I) oxide N2O (dianitrogen monoxide) is obtained from the thermal dissociation of ammonium nitrate:

The molecule has a linear structure

N2O is fairly inert at room temperature, but at high temperatures it can support the combustion of easily oxidized materials. N2O, known as laughing gas, is used for mild anesthesia in medicine. Nitrogen oxide (II) NO is a colorless gas, one of the products of the catalytic thermal dissociation of ammonia in the presence of oxygen:

NO is also formed during the thermal decomposition of nitric acid or during the reaction of copper with dilute nitric acid:

NO can be produced by synthesis from simple substances (N2 and O2) at very high temperatures, for example, in an electrical discharge. The structure of the NO molecule has one unpaired electron. Compounds with this structure interact with electric and magnetic fields. In the liquid or solid state, the oxide is blue in color because the unpaired electron causes partial association in the liquid state and weak dimerization in the solid state: 2NO N2O2. Nitric oxide (III) N2O3 (nitrogen trioxide) - nitrous acid anhydride: N2O3 + H2O 2HNO2. Pure N2O3 can be obtained as a blue liquid at low temperatures (-20° C) from an equimolecular mixture of NO and NO2. N2O3 is stable only in the solid state at low temperatures (melting point -102.3 ° C); in the liquid and gaseous states it again decomposes into NO and NO2. Nitric oxide (IV) NO2 (nitrogen dioxide) also has an unpaired electron in the molecule (see nitric oxide (II) above). The structure of the molecule assumes a three-electron bond, and the molecule exhibits the properties of a free radical (one line corresponds to two paired electrons):

NO2 is obtained by the catalytic oxidation of ammonia in excess oxygen or the oxidation of NO in air:

and also by reactions:

At room temperature, NO2 is a dark brown gas that has magnetic properties due to the presence of an unpaired electron. At temperatures below 0° C, the NO2 molecule dimerizes into dinitrogen tetroxide, and at -9.3° C, dimerization occurs completely: 2NO2 N2O4. In the liquid state, only 1% NO2 is undimerized, and at 100° C 10% N2O4 remains in the form of a dimer. NO2 (or N2O4) reacts in warm water to form nitric acid: 3NO2 + H2O = 2HNO3 + NO. NO2 technology is therefore very important as an intermediate stage in the production of an industrially important product - nitric acid. Nitric oxide (V) N2O5 (obsolete nitric anhydride) is a white crystalline substance obtained by dehydrating nitric acid in the presence of phosphorus oxide P4O10:

N2O5 easily dissolves in air moisture, again forming HNO3. The properties of N2O5 are determined by the equilibrium

N2O5 is a good oxidizing agent; it reacts easily, sometimes violently, with metals and organic compounds and, in its pure state, explodes when heated. The probable structure of N2O5 can be represented as

Nitrogen oxoacids. For nitrogen, three oxoacids are known: hyponitrogenous H2N2O2, nitrogenous HNO2 and nitric acid HNO3. Hyponitrous acid H2N2O2 is a very unstable compound, formed in a non-aqueous medium from a salt of a heavy metal - hyponitrite, under the action of another acid: M2N2O2 + 2HX 2MX + H2N2O2. When the solution is evaporated, a white explosive is formed with the expected structure H-O-N=N-O-H.
Nitrous acid HNO2 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 NO2 (or N2O3) 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. NaNO2 is used in the production of dyes. Nitric acid HNO3 is one of the most important inorganic products of the basic 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.
see also CHEMICAL ELEMENTS.
LITERATURE
Nitrogenist's Handbook. M., 1969 Nekrasov B.V. Fundamentals of general chemistry. M., 1973 Problems of nitrogen fixation. Inorganic and physical chemistry. M., 1982

Collier's Encyclopedia. - Open Society. 2000 .

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