Nitrogen: characteristics, chemical properties, physical properties, compounds, place in nature. Nitrogen gas

Properties elements V-A subgroups

Element	Nitrogen N	Phosphorus R	Arsenic As	Antimony Sb	Bismuth Bi
Property
Element serial number	7	15	33	51	83
Relative atomic mass	14,007	30,974	74,922	121,75	208,980
Melting point, C 0	-210	44,1 (white)	817 (4MPa)	631	271
Boiling point, C 0	-196	280 (white)	613	1380	1560
Density g/cm 3	0,96 (solid)	1,82 (white)	5,72	6,68	9,80
Oxidation states	+5, +3,-3	+5, +3,-3	+5, +3,-3	+5, +3,-3	+5, +3,-3

1. Structure of atoms of chemical elements

Name chemical element	Atomic structure diagram	Electronic structure of the last energy level	Formula of higher oxide R 2 O 5	Volatile hydrogen compound formula RH 3
1. Nitrogen	N+7) 2) 5	…2s 2 2p 3	N2O5	NH 3
2. Phosphorus	P+15) 2) 8) 5	…3s 2 3p 3	P2O5	PH 3
3. Arsenic	As+33) 2) 8) 18) 5	…4s 2 4p 3	As2O5	AsH 3
4. Antimony	Sb+51) 2) 8) 18) 18) 5	…5s 2 5p 3	Sb2O5	SbH 3
5. Bismuth	Bi+83) 2) 8) 18) 32) 18) 5	…6s 2 6p 3	Bi2O5	BiH 3

The presence of three unpaired electrons at the outer energy level explains that in a normal, unexcited state, the valence of elements of the nitrogen subgroup is three.

Atoms of elements of the nitrogen subgroup (except for nitrogen - the outer level of nitrogen consists of only two sublevels - 2s and 2p) have vacant cells of the d-sublevel at the outer energy levels, so they can vaporize one electron from the s-sublevel and transfer it to the d-sublevel . Thus, the valency of phosphorus, arsenic, antimony and bismuth is 5.

Elements of the nitrogen group form compounds of the composition RH 3 with hydrogen, and oxides of the type R 2 O 3 and R 2 O 5 with oxygen. Oxides correspond to acids HRO 2 and HRO 3 (and ortho acids H 3 PO 4, except nitrogen).

The highest oxidation state of these elements is +5, and the lowest is -3.

Since the nuclear charge of atoms increases, the number of electrons per external level constantly, the number of energy levels in atoms increases and the radius of the atom increases from nitrogen to bismuth, the attraction of negative electrons to the positive nucleus weakens and the ability to give up electrons increases, and, consequently, in the nitrogen subgroup, with increasing atomic number, non-metallic properties decrease, and metallic ones increase.

Nitrogen is a non-metal, bismuth is a metal. From nitrogen to bismuth, the strength of RH 3 compounds decreases, and the strength of oxygen compounds increases.

The most important among the elements of the nitrogen subgroup are nitrogen and phosphorus .

Nitrogen, physical and chemical properties, preparation and application

1. Nitrogen is a chemical element

N +7) 2) 5

1 s 2 2 s 2 2 p 3 unfinished external level, p -element, non-metal

Ar(N)=14

2. Possible oxidation states

Due to the presence of three unpaired electrons, nitrogen is very active and is found only in the form of compounds. Nitrogen exhibits oxidation states in compounds from “-3” to “+5”

3. Nitrogen is a simple substance, molecular structure, physical properties

Nitrogen (from Greek ἀ ζωτος - lifeless, lat. Nitrogenium), instead of the previous names (“phlogisticated”, “mephitic” and “spoiled” air) proposed in 1787 Antoine Lavoisier . As shown above, it was already known at that time that nitrogen supports neither combustion nor respiration. This property was considered the most important. Although it later turned out that nitrogen, on the contrary, is essential for all living beings, the name was preserved in French and Russian.

N 2 – covalent nonpolar bond, triple (σ, 2π), molecular crystal lattice

Conclusion:

1. Low reactivity at normal temperature

2. Gas, colorless, odorless, lighter than air

Mr ( B air)/ Mr ( N 2 ) = 29/28

4. Chemical properties of nitrogen

N – oxidizing agent (0 → -3)	N – reducing agent (0 → +5)
1. With metals nitrides are formed Mx Ny - when heated with Mg and alkaline earth and alkaline: 3С a + N 2= Ca 3 N 2 (at t) - c Li at k t room Nitrides are decomposed by water Ca 3 N 2 + 6H 2 O = 3Ca(OH) 2 + 2NH 3 2. With hydrogen 3 H 2 + N 2 ↔ 2 NH 3 (conditions - T, p, kat)	N 2 + O 2 ↔ 2 NO – Q (at t= 2000 C) Nitrogen does not react with sulfur, carbon, phosphorus, silicon and some other non-metals.

5. 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.

In the laboratorypure (“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. EXPERIMENT

6. Application:

In industry, nitrogen gas is used mainly to produce 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; it is used in medicine, especially in cosmetology. Nitrogen mineral fertilizers are important in maintaining soil fertility.

7. Biological role

Nitrogen is an element necessary for the existence of animals and plants; it is part ofproteins (16-18% by weight), amino acids, nucleic acids, nucleoproteins, chlorophyll, hemoglobin etc. In the composition of living cells, the number of nitrogen atoms is about 2%, and the mass fraction is about 2.5% (fourth place after hydrogen, carbon and oxygen). In this regard, a significant amount of fixed nitrogen is contained in living organisms, “dead organic matter” and dispersed matter of the seas and oceans. This amount is estimated at approximately 1.9 10 11 tons. As a result of the processes of rotting and decomposition of nitrogen-containing organic matter, subject to favorable factors environment, natural mineral deposits containing nitrogen may form, for example, the “Chilean saltpeterN 2 → Li 3 N → NH 3

No. 2. Write down equations for the reaction of nitrogen with oxygen, magnesium and hydrogen. For each reaction, create an electronic balance, indicate the oxidizing agent and the reducing agent.

No. 3. One cylinder contains nitrogen gas, another contains oxygen, and the third contains carbon dioxide. How to distinguish these gases?

No. 4. Some flammable gases contain free nitrogen as an impurity. Can the combustion of such gases in ordinary gas stoves nitric oxide (II) is formed. Why?

MOBUSOSH No. 2

Abstract in chemistry on the topic:

“Characteristics of elements of the nitrogen subgroup”

Prepared by: Nasertdinov K.

Checked:

Agidel-2008

2.1.1 Properties of nitrogen

2.1.2 Application of nitrogen

2.2 Ammonia

2.2.1 Properties of ammonia

2.2.2 Application of ammonia

2.2.3 Nitrogen oxides

2.3 Nitric acid

2.3.3 Use of nitric acid and its salts

2.4 Phosphorus

2.4.1 Phosphorus compounds

2.4.2 Application of phosphorus and its compounds

2.5 Mineral fertilizers

Literature

1. Characteristics of elements of the nitrogen subgroup

Nitrogen is the most important component atmosphere (78% of its volume). In nature, it is found in proteins, in deposits of sodium nitrate. Natural nitrogen consists of two isotopes: 14 N (99.635% mass) and 15 N (0.365% mass).

Phosphorus is part of all living organisms. Occurs in nature in the form of minerals. Phosphorus is widely used in medicine, agriculture, aviation, in the extraction of precious metals.

Arsenic, antimony and bismuth are quite widespread, mainly in the form of sulfide ores. Arsenic is one of the elements of life that promotes hair growth. Arsenic compounds are poisonous, but in small doses they can have medicinal properties. Arsenic is used in medicine and veterinary medicine.

2. Structure and characteristics of atoms

Subgroup elements on the outer electrolayer have five electrons. They can give them away, and they can attract three more electrons from other atoms. Therefore, their oxidation state is from -3 to +5. Their volatile hydrogen and higher oxygen compounds are acidic in nature and are designated general formulas: RH 3 and R 2 O 5.

The elements of the subgroup have non-metallic properties, and at the same time the ability to attract electrons is less than that of the elements of the halogen and oxygen subgroups.

In the nitrogen subgroup in the periodic table, as elements move from top to bottom, metallic properties increase.

Nitrogen and phosphorus are non-metals, arsenic and antimony exhibit properties of metals, and bismuth is a metal.

Substance name	Molecular formula	Structure	Physical properties	Density, g/cm 3	Temperature, about C
Nitrogen	N 2	Molecular	Gas without color, smell, taste, soluble in water	0.81 (w)	plv		bale
					-210		-195,8
Phosphorus white	P 4	Tetrahedral molecule. Molecular crystal lattice.	Soft solid, colorless, slightly soluble in water, soluble in carbon sulfur	1,82	44 (underwater)		257
Arsenic gray	As 4	Same.	Brittle crystalline substance with metal. shine on a fresh break. Insoluble in water. Very weak conductor of electricity	5,72	Sublimates, passes from solid to gaseous (steam) at 615 o C
Antimony	Sb 4	--	Silvery-white crystalline substance, brittle, poor conductor of heat and electricity	6,68	630,5	1634
Bismuth	Bi n	A molecular crystal in which each atom is bonded to three neighboring ones.	Pink-white, brittle crystalline substance, resembling metal in appearance, electrical conductivity is negligible	9,8	271,3	1550

Table of properties of simple substances of elements of the nitrogen subgroup.

2.1 Nitrogen

Nitrogen is the initial and most important element of the subgroup. Nitrogen is a typical non-metallic element. Unlike other elements of the subgroup, nitrogen does not have the ability to increase its valency. The electronic structure is represented by seven electrons located at two energy levels. Electronic formula: 1s 2 2s 2 2p 3. Oxidation degrees of nitrogen: - 3,+5,-2,-1,+1,+2,+3,+4. The nitrogen atom has high chemical activity; it attaches electrons more actively than sulfur and phosphorus atoms.

2.1.1 Properties of nitrogen

Nitrogen at normal conditions- molecular, gaseous, low-active substance, the molecule consists of two atoms; colorless gas, odorless, slightly soluble in water, slightly lighter than air, does not react with oxygen, at - 196 o C it compresses, at - 210 o C it turns into a snow-like mass.

Nitrogen is chemically inactive. It does not support either breathing or combustion. At room temperature, it reacts only with lithium, forming Li 3 N. To break a nitrogen molecule, 942 kJ/mol of energy must be expended. The reactions in which nitrogen enters are redox, where nitrogen exhibits the properties of both an oxidizing agent and a reducing agent.

At elevated temperatures, nitrogen combines with many metals, at room temperature - only with lithium. Nitrogen interacts with non-metals at an even higher temperature. Thanks to this, life on our planet is possible, since if nitrogen reacted at low temperatures, it would react with oxygen, which is part of the air, and living beings would not be able to breathe this mixture of gases.

2.1.2 Application of nitrogen

Nitrogen in industry is obtained from the air using the difference in boiling points of nitrogen and oxygen.

Nitrogen is used in the chemical industry to produce ammonia, urea, etc.; in electrical engineering when creating electric lamps, pumping flammable liquids, drying explosives, etc.

2.2 Ammonia

Ammonia is one of the most important hydrogen compounds of nitrogen. It is of great practical importance. Life on Earth owes much to certain bacteria that can convert atmospheric nitrogen into ammonia.

2.2.1 Properties of ammonia

The ammonia molecule is formed by pairing three p-electrons of a nitrogen atom with three s-electrons of hydrogen atoms. Oxidation state: - 3. The ammonia molecule is highly polar.

Ammonia is a colorless gas with a pungent odor, almost twice as light as air. When cooled to - 33 o C, it contracts. Ammonia is highly soluble in water.

Ammonia is a chemically active compound that reacts with many substances. Most often these are oxidation and compound reactions. In redox reactions, ammonia acts only as a reducing agent. Ammonia burns in oxygen and combines actively with water and acids.

2.2.2 Application of ammonia

Ammonia is used to produce nitric acid and nitrogen-containing mineral fertilizers, salts, soda. In liquid form, it is used in refrigeration. Ammonia is used in medicine to create ammonia; in everyday life as part of stain removers, as well as in chemical laboratories. Ammonium salts are used for the production of explosives, fertilizers, electric batteries, and for metal processing and welding.

2.2.3 Nitrogen oxides

For nitrogen, oxides are known that correspond to all its positive oxidation states (+1,+2,+3,+4,+5): N 2 O, NO, N 2 O 3, NO 2, N 2 O 4, N 2 O 5 . Under normal conditions, nitrogen does not interact with oxygen, only when an electric discharge is passed through their mixture.

Table of properties of nitrogen oxides.

2.3 Nitric acid

2.3.1 Properties of nitric acid

The nitric acid molecule HNO 3 consists of three elements connected to each other by covalent bonds. This is a molecular substance containing a highly oxidized nitrogen atom. However, the valence of nitrogen in the acid is four instead of the usual oxidation number of nitrogen.

Pure nitric acid is a colorless liquid, fuming in air, with a pungent odor. Concentrated nitric acid is colored yellow. The density of nitric acid is 1.51 g/cm 3, the boiling point is 86 o C, and at a temperature of 41.6 o C it solidifies in the form of a transparent crystalline mass. The acid dissolves in water and the aqueous solution is an electrolyte.

Dilute nitric acid exhibits properties common to all acids. It is a strong oxidizing agent. At room temperature, the acid decomposes into nitric oxide (IV), oxygen and water, so it is stored in dark bottles in a cool place. It reacts with metals (except gold and platinum), both active and inactive.

Many nonmetals are oxidized by nitric acid. Nitric acid, especially concentrated acid, oxidizes organic substances. Animal and plant tissues are quickly destroyed when exposed to nitric acid.

2.3.2 Salts of nitric acid and their properties

Salts of nitric acid, nitrates, are formed when the acid reacts with metals, metal oxides, bases, ammonia, and also with some salts.

Nitrates are crystalline solids that dissolve well in water. strong electrolytes. When heated, they decompose releasing oxygen. It has a number of specific properties as an oxidizing agent. Depending on the nature of the metal, the decomposition reaction proceeds differently.

A qualitative reaction to nitrate ion (solutions of nitric acid and its salt) is carried out as follows: copper shavings are added to the test tube with the test substance, sulfuric acid concentrate is added and heated. The release of brown gas indicates the presence of nitrate ion.

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

Distribution of Nitrogen in nature. Nitrogen is one of the most common elements on Earth, and the bulk of it (about 4·10 15 tons) is concentrated in a free state in the atmosphere. In the air, free Nitrogen (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 nitrogen content in the lithosphere is 1.9·10 -3% by mass. Natural nitrogen compounds are 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 Nitrogen for industry (now the industrial synthesis of ammonia from Nitrogen in air and hydrogen is of primary importance for fixing Nitrogen). Small amounts of bound Nitrogen are found in coal (1-2.5%) and oil (0.02-1.5%), as well as in the waters of rivers, seas and oceans. Nitrogen accumulates in soils (0.1%) and in living organisms (0.3%).

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

In nature, the Nitrogen cycle takes place, in which the main role is played by microorganisms - nitrophying, denitrophying, nitrogen-fixing and others. However, as a result of the extraction of huge amounts of bound Nitrogen from the soil by plants (especially during intensive farming), the soils become depleted of Nitrogen. Nitrogen deficiency is typical for agriculture in almost all countries; nitrogen deficiency is also observed in animal husbandry (“protein starvation”). On soils poor in available Nitrogen, plants develop poorly. Nitrogen fertilizers and protein feeding of animals are the most important means of boosting agriculture. Economic activity humans disrupt the nitrogen cycle. Thus, burning fuel enriches the atmosphere with Nitrogen, and factories producing fertilizers bind Nitrogen from the air. Transportation of fertilizers and agricultural products redistributes Nitrogen to the surface of the earth. Nitrogen is the fourth most abundant element in the solar system (after hydrogen, helium and oxygen).

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

Physical properties of Nitrogen. Nitrogen is slightly lighter than air; density 1.2506 kg/m 3 (at 0°C and 101325 n/m 2 or 760 mm Hg), melting point -209.86°C, boiling point -195.8°C. Nitrogen 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); the density of liquid nitrogen is 808 kg/m3. In water, Nitrogen is less soluble than oxygen: at 0°C, 23.3 g of Nitrogen dissolves in 1 m 3 H 2 O. Nitrogen is soluble in some hydrocarbons better than in water.

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

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

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

Obtaining Nitrogen. In the laboratory, Nitrogen can be easily obtained by heating a concentrated solution of ammonium nitrite: NH 4 NO 2 = N 2 + 2H 2 O. Technical method Nitrogen production is based on the separation of pre-liquefied air, which is then subjected to distillation.

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

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

Nitrogen in the body. Nitrogen is one of the main biogenic elements that make up the most important substances of living cells - proteins and nucleic acids. However, the amount of Nitrogen in the body is small (1-3% by dry weight). Only some microorganisms and blue-green algae can assimilate molecular nitrogen in the atmosphere.

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

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

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

Nitrogen- element of the 2nd period of the V A-group of the Periodic Table, serial number 7. Electronic formula of the atom [ 2 He]2s 2 2p 3, characteristic oxidation states 0, -3, +3 and +5, less often +2 and +4, etc. the N v state is considered relatively stable.

Scale of oxidation states for nitrogen:
+5 - N 2 O 5, NO 3, NaNO 3, AgNO 3

3 – N 2 O 3, NO 2, HNO 2, NaNO 2, NF 3

3 - NH 3, NH 4, NH 3 * H 2 O, NH 2 Cl, Li 3 N, Cl 3 N.

Nitrogen has a high electronegativity (3.07), third after F and O. It exhibits typical non-metallic (acidic) properties, forming various oxygen-containing acids, salts and binary compounds, as well as the ammonium cation NH 4 and its salts.

In nature - seventeenth by chemical abundance element (ninth among non-metals). A vital element for all organisms.

N 2

Simple substance. It consists of non-polar molecules with a very stable ˚σππ-bond N≡N, this explains the chemical inertness of the element under normal conditions.

A colorless, tasteless and odorless gas that condenses into a colorless liquid (unlike O2).

The main component of air is 78.09% by volume, 75.52 by mass. Nitrogen boils away from liquid air before oxygen does. Slightly soluble in water (15.4 ml/1 l H 2 O at 20 ˚C), the solubility of nitrogen is less than that of oxygen.

At room temperature N2 reacts with fluorine and, to a very small extent, with oxygen:

N 2 + 3F 2 = 2NF 3, N 2 + O 2 ↔ 2NO

The reversible reaction to produce ammonia occurs at a temperature of 200˚C, under pressure up to 350 atm and always in the presence of a catalyst (Fe, F 2 O 3, FeO, in the laboratory with Pt)

N 2 + 3H 2 ↔ 2NH 3 + 92 kJ

According to Le Chatelier's principle, an increase in ammonia yield should occur with increasing pressure and decreasing temperature. However, the reaction rate at low temperatures is very small, so the process is carried out at 450-500 ˚C, achieving a 15% ammonia yield. Unreacted N 2 and H 2 are returned to the reactor and thereby increase the degree of reaction.

Nitrogen is chemically passive in relation to acids and alkalis and does not support combustion.

Receipt V industry– fractional distillation of liquid air or removal of oxygen from air by chemical means, for example, by the reaction 2C (coke) + O 2 = 2CO when heated. In these cases, nitrogen is obtained, which also contains impurities of noble gases (mainly argon).

In the laboratory, small amounts of chemically pure nitrogen can be obtained by the commutation reaction with moderate heating:

N -3 H 4 N 3 O 2(T) = N 2 0 + 2H 2 O (60-70)

NH 4 Cl(p) + KNO 2 (p) = N 2 0 + KCl + 2H 2 O (100˚C)

Used for ammonia synthesis. Nitric acid and other nitrogen-containing products, as an inert medium for chemical and metallurgical processes and storage of flammable substances.

N.H. 3

Binary compound, the oxidation state of nitrogen is – 3. Colorless gas with a sharp characteristic odor. The molecule has the structure of an incomplete tetrahedron [: N(H) 3 ] (sp 3 hybridization). The presence of a donor pair of electrons on the sp 3 hybrid orbital of nitrogen in the NH 3 molecule determines the characteristic reaction of addition of a hydrogen cation, which results in the formation of a cation ammonium NH4. It liquefies under excess pressure at room temperature. In the liquid state, it is associated through hydrogen bonds. Thermally unstable. Highly soluble in water (more than 700 l/1 l H 2 O at 20˚C); the share in a saturated solution is 34% by weight and 99% by volume, pH = 11.8.

Very reactive, prone to addition reactions. Burns in oxygen, reacts with acids. It exhibits reducing (due to N -3) and oxidizing (due to H +1) properties. It is dried only with calcium oxide.

Qualitative reactions – the formation of white “smoke” upon contact with gaseous HCl, blackening of a piece of paper moistened with a solution of Hg 2 (NO3) 2.

An intermediate product in the synthesis of HNO 3 and ammonium salts. Used in the production of soda, nitrogen fertilizers, dyes, explosives; liquid ammonia is a refrigerant. Poisonous.
Equations of the most important reactions:

2NH 3 (g) ↔ N 2 + 3H 2
NH 3 (g) + H 2 O ↔ NH 3 * H 2 O (p) ↔ NH 4 + + OH —
NH 3 (g) + HCl (g) ↔ NH 4 Cl (g) white “smoke”
4NH 3 + 3O 2 (air) = 2N 2 + 6 H 2 O (combustion)
4NH 3 + 5O 2 = 4NO+ 6 H 2 O (800˚C, cat. Pt/Rh)
2 NH 3 + 3CuO = 3Cu + N 2 + 3 H 2 O (500˚C)
2 NH 3 + 3Mg = Mg 3 N 2 +3 H 2 (600 ˚C)
NH 3 (g) + CO 2 (g) + H 2 O = NH 4 HCO 3 (room temperature, pressure)
Receipt. IN laboratories– displacement of ammonia from ammonium salts when heated with soda lime: Ca(OH) 2 + 2NH 4 Cl = CaCl 2 + 2H 2 O + NH 3
Or boiling an aqueous solution of ammonia and then drying the gas.
In industry Ammonia is produced from nitrogen and hydrogen. Produced by industry either in liquefied form or in the form of a concentrated aqueous solution under the technical name ammonia water.

Ammonia hydrateN.H. 3 * H 2 O. Intermolecular connection. White, in the crystal lattice – NH 3 and H 2 O molecules connected by a weak hydrogen bond. Present in an aqueous solution of ammonia, a weak base (dissociation products - NH 4 cation and OH anion). The ammonium cation has a regular tetrahedral structure (sp 3 hybridization). Thermally unstable, completely decomposes when the solution is boiled. Neutralized by strong acids. Shows reducing properties (due to N-3) in a concentrated solution. It undergoes ion exchange and complexation reactions.

Qualitative reaction– formation of white “smoke” upon contact with gaseous HCl. It is used to create a slightly alkaline environment in solution during the precipitation of amphoteric hydroxides.
A 1 M ammonia solution contains mainly NH 3 *H 2 O hydrate and only 0.4% NH 4 OH ions (due to hydrate dissociation); Thus, the ionic “ammonium hydroxide NH 4 OH” is practically not contained in the solution, and there is no such compound in the solid hydrate.
Equations of the most important reactions:
NH 3 H 2 O (conc.) = NH 3 + H 2 O (boiling with NaOH)
NH 3 H 2 O + HCl (diluted) = NH 4 Cl + H 2 O
3(NH 3 H 2 O) (conc.) + CrCl 3 = Cr(OH) 3 ↓ + 3 NH 4 Cl
8(NH 3 H 2 O) (conc.) + 3Br 2(p) = N 2 + 6 NH 4 Br + 8H 2 O (40-50˚C)
2(NH 3 H 2 O) (conc.) + 2KMnO 4 = N 2 + 2MnO 2 ↓ + 4H 2 O + 2KOH
4(NH 3 H 2 O) (conc.) + Ag 2 O = 2OH + 3H 2 O
4(NH 3 H 2 O) (conc.) + Cu(OH) 2 + (OH) 2 + 4H 2 O
6(NH 3 H 2 O) (conc.) + NiCl 2 = Cl 2 + 6H 2 O
A dilute ammonia solution (3-10%) is often called ammonia(the name was invented by alchemists), and the concentrated solution (18.5 - 25%) is an ammonia solution (produced by industry).

Nitrogen oxides

Nitrogen monoxideNO

Non-salt-forming oxide. Colorless gas. Radical, contains a covalent σπ bond (N꞊O), in the solid state a dimer of N 2 O 2 co N-N connection. Extremely thermally stable. Sensitive to air oxygen (turns brown). Slightly soluble in water and does not react with it. Chemically passive towards acids and alkalis. When heated, it reacts with metals and non-metals. a highly reactive mixture of NO and NO 2 (“nitrous gases”). Intermediate product in the synthesis of nitric acid.
Equations of the most important reactions:
2NO + O 2 (g) = 2NO 2 (20˚C)
2NO + C (graphite) = N 2 + CO 2 (400-500˚C)
10NO + 4P(red) = 5N 2 + 2P 2 O 5 (150-200˚C)
2NO + 4Cu = N 2 + 2 Cu 2 O (500-600˚C)
Reactions to mixtures of NO and NO 2:
NO + NO 2 +H 2 O = 2HNO 2 (p)
NO + NO 2 + 2KOH(dil.) = 2KNO 2 + H 2 O
NO + NO 2 + Na 2 CO 3 = 2Na 2 NO 2 + CO 2 (450-500˚C)
Receipt V industry: oxidation of ammonia with oxygen on a catalyst, in laboratories— interaction of dilute nitric acid with reducing agents:
8HNO 3 + 6Hg = 3Hg 2 (NO 3) 2 + 2 NO+ 4 H 2 O
or nitrate reduction:
2NaNO 2 + 2H 2 SO 4 + 2NaI = 2 NO + I 2 ↓ + 2 H 2 O + 2Na 2 SO 4

Nitrogen dioxideNO 2

Acid oxide, conditionally corresponds to two acids - HNO 2 and HNO 3 (acid for N 4 does not exist). Brown gas, at room temperature a monomer NO 2, in the cold a liquid colorless dimer N 2 O 4 (dianitrogen tetroxide). Reacts completely with water and alkalis. A very strong oxidizing agent that causes corrosion of metals. It is used for the synthesis of nitric acid and anhydrous nitrates, as a rocket fuel oxidizer, an oil purifier from sulfur, and a catalyst for the oxidation of organic compounds. Poisonous.
Equation of the most important reactions:
2NO 2 ↔ 2NO + O 2
4NO 2 (l) + H 2 O = 2HNO 3 + N 2 O 3 (syn.) (in the cold)
3 NO 2 + H 2 O = 3HNO 3 + NO
2NO 2 + 2NaOH (diluted) = NaNO 2 + NaNO 3 + H 2 O
4NO 2 + O 2 + 2 H 2 O = 4 HNO 3
4NO 2 + O 2 + KOH = KNO 3 + 2 H 2 O
2NO 2 + 7H 2 = 2NH 3 + 4 H 2 O (cat. Pt, Ni)
NO 2 + 2HI(p) = NO + I 2 ↓ + H 2 O
NO 2 + H 2 O + SO 2 = H 2 SO 4 + NO (50-60˚C)
NO 2 + K = KNO 2
6NO 2 + Bi(NO 3) 3 + 3NO (70-110˚C)
Receipt: V industry - oxidation of NO by atmospheric oxygen, in laboratories– interaction of concentrated nitric acid with reducing agents:
6HNO 3 (conc., hor.) + S = H 2 SO 4 + 6NO 2 + 2H 2 O
5HNO 3 (conc., hor.) + P (red) = H 3 PO 4 + 5NO 2 + H 2 O
2HNO 3 (conc., hor.) + SO 2 = H 2 SO 4 + 2 NO 2

Dianitrogen oxideN 2 O

A colorless gas with a pleasant odor (“laughing gas”), N꞊N꞊О, formal oxidation state of nitrogen +1, poorly soluble in water. Supports combustion of graphite and magnesium:

2N 2 O + C = CO 2 + 2N 2 (450˚C)
N 2 O + Mg = N 2 + MgO (500˚C)
Obtained by thermal decomposition of ammonium nitrate:
NH 4 NO 3 = N 2 O + 2 H 2 O (195-245˚C)
used in medicine as an anesthetic.

Dianitrogen trioxideN 2 O 3

At low temperatures – blue liquid, ON꞊NO 2, formal oxidation state of nitrogen +3. At 20 ˚C, it decomposes 90% into a mixture of colorless NO and brown NO 2 (“nitrous gases”, industrial smoke – “fox tail”). N 2 O 3 – acid oxide, in the cold with water forms HNO 2, when heated it reacts differently:
3N 2 O 3 + H 2 O = 2HNO 3 + 4NO
With alkalis it gives salts HNO 2, for example NaNO 2.
Obtained by reacting NO with O 2 (4NO + 3O 2 = 2N 2 O 3) or with NO 2 (NO 2 + NO = N 2 O 3)
with strong cooling. “Nitrous gases” are also environmentally dangerous and act as catalysts for the destruction of the ozone layer of the atmosphere.

Dianitrogen pentoxide N 2 O 5

Colorless, solid substance, O 2 N – O – NO 2, nitrogen oxidation state is +5. At room temperature it decomposes into NO 2 and O 2 in 10 hours. Reacts with water and alkalis as an acid oxide:
N2O5 + H2O = 2HNO3
N 2 O 5 + 2NaOH = 2NaNO 3 + H 2
Prepared by dehydration of fuming nitric acid:
2HNO3 + P2O5 = N2O5 + 2HPO3
or oxidation of NO 2 with ozone at -78˚C:
2NO 2 + O 3 = N 2 O 5 + O 2

Nitrites and nitrates

Potassium nitriteKNO 2 . White, hygroscopic. Melts without decomposition. Stable in dry air. Very soluble in water (forming a colorless solution), hydrolyzes at the anion. A typical oxidizing and reducing agent in an acidic environment, it reacts very slowly in an alkaline environment. Enters into ion exchange reactions. Qualitative reactions on the NO 2 ion - discoloration of the violet MnO 4 solution and the appearance of a black precipitate when adding I ions. It is used in the production of dyes, as an analytical reagent for amino acids and iodides, and a component of photographic reagents.
equation of the most important reactions:
2KNO 2 (t) + 2HNO 3 (conc.) = NO 2 + NO + H 2 O + 2KNO 3
2KNO 2 (dil.)+ O 2 (e.g.) → 2KNO 3 (60-80 ˚C)
KNO 2 + H 2 O + Br 2 = KNO 3 + 2HBr
5NO 2 - + 6H + + 2MnO 4 - (viol.) = 5NO 3 - + 2Mn 2+ (bts.) + 3H 2 O
3 NO 2 - + 8H + + CrO 7 2- = 3NO 3 - + 2Cr 3+ + 4H 2 O
NO 2 - (saturated) + NH 4 + (saturated) = N 2 + 2H 2 O
2NO 2 - + 4H + + 2I - (bts.) = 2NO + I 2 (black) ↓ = 2H 2 O
NO 2 - (diluted) + Ag + = AgNO 2 (light yellow)↓
Receipt Vindustry– reduction of potassium nitrate in the processes:
KNO3 + Pb = KNO 2+ PbO (350-400˚C)
KNO 3 (conc.) + Pb (sponge) + H 2 O = KNO 2+ Pb(OH) 2 ↓
3 KNO3 + CaO + SO2 = 2 KNO 2+ CaSO 4 (300 ˚C)

H itrate potassium KNO 3
Technical name potash, or Indian salt , saltpeter. White, melts without decomposition and decomposes upon further heating. Stable in air. Highly soluble in water (with high endo-effect, = -36 kJ), no hydrolysis. A strong oxidizing agent during fusion (due to the release of atomic oxygen). In solution it is reduced only by atomic hydrogen (in an acidic environment to KNO 2, in an alkaline environment to NH 3). Used in glass production as a preservative food products, a component of pyrotechnic mixtures and mineral fertilizers.

2KNO 3 = 2KNO 2 + O 2 (400-500 ˚C)

KNO 3 + 2H 0 (Zn, dil. HCl) = KNO 2 + H 2 O

KNO 3 + 8H 0 (Al, conc. KOH) = NH 3 + 2H 2 O + KOH (80 ˚C)

KNO 3 + NH 4 Cl = N 2 O + 2H 2 O + KCl (230-300 ˚C)

2 KNO 3 + 3C (graphite) + S = N 2 + 3CO 2 + K 2 S (combustion)

KNO 3 + Pb = KNO 2 + PbO (350 - 400 ˚C)

KNO 3 + 2KOH + MnO 2 = K 2 MnO 4 + KNO 2 + H 2 O (350 - 400 ˚C)

Receipt: in industry
4KOH (hor.) + 4NO 2 + O 2 = 4KNO 3 + 2H 2 O

and in the laboratory:
KCl + AgNO 3 = KNO 3 + AgCl↓

The content of the article

NITROGEN, N (nitrogenium), chemical element (at. number 7) VA subgroup periodic table 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 represents a small fraction of the nitrogen contained in 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. Vegetable and animal world contains 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 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.” Latin name 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 leguminous plants, for example, peas, clover and soybeans, accumulate nodules on their roots, in which nitrogen-fixing bacteria convert it into nitrates, or oxidation of atmospheric nitrogen with oxygen occurs under lightning discharge 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: c mass number 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 . Therefore, on the outer (second) electron shell there are 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 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 different ways, 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 at all, 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 when 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 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 type 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.

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-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. Main industrial method The production of 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.

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 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. Aqueous ammonia solution (" ammonia") consists predominantly of NH 3, H 2 O and low 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 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 oxides heavy metals to metals at high temperatures, apparently due to hydrogen formed as a result of 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 final 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 salts hydrazonium type + X – . 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 with the formation of 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



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