Scheme of the structure of the nitrogen atom n 20. Electronic structure of the pyrrole nitrogen atom. Chemical properties of nitric acid

In the pyridine molecule, p,p conjugation takes place. Pyridinic nitrogen, due to its greater electronegativity compared to carbon, shifts the single p-electron density towards itself, generally reducing the electron density of the aromatic ring. Therefore, such systems with pyridine nitrogen are called p-deficient.

When replacing the fragment - CH = CH - with > NH, a five-membered ring appears - pyrrole

1. The pyrrole molecule has a cyclic structure.

2. All carbon atoms in the cycle are in sp 2 hybridization, the nitrogen atom is also sp 2 hybridized, and the nitrogen atom supplies a two-electron P z orbital to a single p-electron cloud.

3. The total π electron density of pyrrole includes 4n+2 = 6 p electrons

In the pyrrole molecule, p,p conjugation takes place. Systems containing pyrrole nitrogen are called p-excess or superaromatic systems. The presence of such a system greatly affects the reactivity of pyrrole.

In natural compounds, the aromatic pyrrole ring is often found in various polynuclear compounds, of which the most important is the porphin nucleus, which is part of hemoglobin and chlorophyll.

A conjugated system of 26 p-electrons (11 double bonds and 2 lone pairs of electrons of pyrrole atoms. The high conjugation energy (840 KJ) indicates the high stability of porphine.

The concept of aromaticity extends not only to neutral molecules, but also to charged ions. _

When replacing the fragment – ​​CH=CH – in benzene with – CH, a carbocyclic – cyclopentadienyl anion arises, which belongs to the non-benzenoid structure. Cyclopentadienyl ion is part of medicinal product ferrocene (dicyclopentadienyl iron) and the natural compound azulene.

The cyclopentadienyl anion is formed by the abstraction of a proton from cyclopentadiene-1,3.

Let's consider the aromaticity criteria for the cyclopentadienyl anion:

1) cyclic connection

2) all carbon atoms have sp 2 hybridization

Ferrocene is a sandwich-like organometallic compound (stimulates hematopoiesis and is used for iron deficiency anemia.

The cycloheptatrienyl cation (tropylium cation) is formed from cycloheptatriene-1,3,5 by elimination of the hydride ion.

The tropylium cation is a regular heptagon. An aromatic sextet is formed by the overlap of 6 one-electron and one vacant p z orbital.

Let's consider the aromaticity criteria for tropylium cation:

1) The connection is cyclic

2) All carbon atoms have sp 2 hybridization

3) General π -electronic system includes 4n + 2 = 6 p electrons

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. Chemical and physical properties and these substances, the structure of their molecules, reactions with other substances. In addition, methods for obtaining these substances by industrial and laboratory methods and their use in various industries will be listed. Review the properties and uses of nitrous oxide and aqua regia(a compound 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.

Examples of substances with varying degrees oxidation 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.

The element nitrogen N is the first representative of the main subgroup of group V of the Periodic table. Its atoms contain five electrons in the outer energy level, of which three are unpaired electrons (remember the “8-N” rule). It follows that the atoms of these elements can add three electrons, completing the external energy level, and as a result acquire an oxidation state of -3, for example, in compounds with hydrogen - ammonia NH3 and with metals - nitrides Li3N, Mg3N2, etc.
Nitrogen atoms can also donate their outer electrons more electronegative elements (fluorine, oxygen) and acquire oxidation states +3 and +5. Nitrogen atoms also exhibit reducing properties in oxidation states +1, +2, +4.

Nitrogen - simple substance. In the free state, nitrogen exists in the form of a diatomic molecule M2. In this molecule, two N atoms are connected by a very strong triple covalent bond:

This connection can also be expressed as follows:
N=N

Nitrogen is a colorless, odorless and tasteless gas. It is less soluble in water than oxygen. The strength of the nitrogen molecule is due to its chemical inertness.
Under normal conditions, nitrogen reacts only with lithium, forming nitride Li3N:
6Li + N2 = 2Li3N
It interacts with other metals only when high temperatures.
Also, at high temperatures and pressures in the presence of a catalyst, nitrogen reacts with hydrogen to form ammonia:
2N+ ZN2<->2NH3
(characterize this reaction and consider the conditions for the shift of chemical equilibrium to the right).
At the temperature of the electric arc, it combines with oxygen, forming nitrogen oxide (II) (characterize this reaction and also consider the conditions for the shift of chemical equilibrium to the right).
In nature, nitrogen is found mainly in the atmosphere - 78.09% by volume or 65.6% by mass. 8 thousand tons of nitrogen are constantly “hanging” over each hectare of the earth’s surface. Of the natural inorganic nitrogen compounds, the most famous is Chilean saltpeter NaNO3.
Most of the fixed nitrogen is found in organic matter.
Nitrogen obtained by distillation of liquid air is used in industry for the synthesis of ammonia and the production of nitric acid. Previously, this gas was used as an inert medium to fill electric lamps. In medicine, pure nitrogen is used as an inert medium in the treatment of pulmonary tuberculosis, and a liquid nitrogen- in the treatment of diseases of the spine, joints, etc.

Nitrogen cycle in nature. Nitrogen is a vital element. All the main parts of the cells of the body's tissues are built from protein molecules, which include nitrogen. Without protein there is no life, and without nitrogen there is no protein. Man gets proteins from plant and animal foods, and animals, in turn, also get them from plants. Therefore, plants are one of the sources of nitrogen replenishment that supports life.
The content of fixed nitrogen in the soil is very small (up to 1 kg per 1 ton), moreover, most of it is part of organic compounds and is not directly available to plants. However, gradually, as a result of the activity of bacteria, organic compounds are converted into mineral compounds - ammonium salts or nitrates, which are absorbed by plants.
Nitrogen is part of plant proteins. Animals obtain ready-made protein substances from plants; The animal body contains from 1 to 10% nitrogen (by weight), wool and horns contain about 15%. All the most important parts of cells (cytoplasm, nucleus, membrane) are built from protein molecules.
Of even greater importance are special bacteria that live in nodules on the roots of leguminous plants (clover, peas, vetch, lupine, etc.); they are called “nodule bacteria”. It is these bacteria that bind free atmospheric nitrogen, that is, they convert it into compounds that plants absorb, forming proteins in their body.
Nitrogen compounds in the soil are also replenished during thunderstorms. As you already know, in this case, nitric oxide (N) is formed from nitrogen and oxygen, which, under the influence of atmospheric oxygen, turns into nitric oxide (IV):
2NO + 02 = 2NO2
The latter reacts with water (also in the presence of atmospheric oxygen), and nitric acid is obtained:
4NO2 + 02 + 2H20 = 4HNO3

This acid, entering the soil, reacts with the compounds of sodium, calcium, and potassium found in it and forms salts - nitrate, necessary for plants (Fig. 27).
Discovery of nitrogen . In 1772, the English scientist D. Rutherford and the Swedish researcher K. Scheele discovered in their experiments on the combustion of substances a gas that does not support respiration and combustion. Later, in 1787, A. Lavoisier established the presence in the air of a gas that does not support respiration and combustion, and at his suggestion this gas was given the name “nitrogen”, meaning “lifeless” (from the Latin a - no and zoe - life ). Another Latin name nitrogenium, introduced in 1790 by J. Chaptal, means “giving birth to saltpeter.”

Ammonia

First of all, let's look at the structure of the ammonia molecule NH3. As you already know, at the outer energy level, nitrogen atoms contain five electrons, of which three are unpaired electrons. It is they who participate in the formation of three covalent bonds with three hydrogen atoms during the formation of the ammonia molecule NH3:

Three common electron pairs are shifted towards the more electronegative nitrogen atom, and since the ammonia molecule has the shape of a triangular pyramid (Fig. 28), as a result of the displacement of electron pairs, a dipole appears, i.e. a system with two poles.

Hydrogen bond- This is a chemical bond between the hydrogen atoms of one molecule and the atoms of very electronegative elements (fluorine, oxygen, nitrogen) that have shared electron pairs of another molecule.
This is a very weak bond - about 15-20 times weaker than a covalent bond. Thanks to it, some low-molecular substances (i.e., having a small molecular weight) form associates, which leads to an increase in the melting and boiling points of substances. Hydrogen bonds are formed between molecules of water, alcohols, and hydrogen fluoride.
Hydrogen bonds play a very important role in the molecules of the most important compounds for living beings - proteins and nucleic acids.
Ammonia - a colorless gas with a pungent odor, almost twice as light as air. Ammonia should not be inhaled for long periods of time as it is poisonous. This gas easily liquefies at normal pressure and a temperature of -33.4 ° C, and when liquid ammonia evaporates from the environment, a lot of heat is absorbed, which is why ammonia is used in refrigeration units.
Ammonia is very soluble in water: at 20 °C, about 710 volumes of ammonia are dissolved in 1 volume of water (Fig. 29). A concentrated aqueous solution of ammonia (25% by weight) is called aqueous ammonia, or ammonia water, and the ammonia solution used in medicine is known as ammonia. The ammonia found in your home medicine cabinet contains 10% ammonia.
If you add a few drops of phenolphthalein to an ammonia solution, it will turn crimson, i.e. it will show an alkaline environment:
NH3 + H20<->NH3 H20 -> NH4 + OH-
The presence of hydroxide ions OH- explains the alkaline reaction of aqueous ammonia solutions. If an ammonia solution colored with phenolphthalein is heated, the color will disappear (why?).

Ammonia reacts with acids to form ammonium salts. This interaction is clearly observed in the following experiment: if a glass rod or glass moistened with an ammonia solution is brought to another rod or glass moistened with a solution of hydrochloric acid, thick white smoke will appear (Fig. 30). So believe after this saying that there is no smoke without fire:
NH3 + HCl = NH4Сl
Ammonium chloride
Both an aqueous solution of ammonia and ammonium salts contain a special ion - the ammonium cation NH4, which plays the role of a metal cation. It is obtained as a result of the fact that the nitrogen atom has a free (lone) electron pair, due to which another covalent bond is formed with the hydrogen cation, which is transferred to ammonia from acid or water molecules:

This mechanism for the formation of a covalent bond, which arises not as a result of the sharing of unpaired electrons, but due to a free electron pair present in one of the atoms, is called donor-acceptor.

In this case, the donor of this free electron pair is the nitrogen atom in ammonia, and the acceptor is the hydrogen cation of acid or water.
You can predict another chemical property of ammonia yourself if you pay attention to the oxidation state of nitrogen atoms in it, namely -3. Of course, ammonia is the strongest reducing agent, that is, its nitrogen atoms can only give up electrons, but not accept them. Thus, ammonia can be oxidized either to free nitrogen (without the participation of a catalyst):
4NН3 + 302 = 2N2 + 6Н20
or to nitrogen oxide(II) (in the presence of a catalyst):
4NН3 + 502 = 4N + 6Н20
You already know how ammonia is produced in industry - by synthesis from nitrogen and hydrogen. In the laboratory, ammonia is obtained by the action of slaked lime Ca(OH)2 on ammonium salts, most often ammonium chloride:
Ca(OH)2 + 2NH4C1 = CaCl2 + 2NH3 + 2H20

The gas is collected in a vessel turned upside down, and is recognized either by smell, or by the blueness of wet red litmus paper, or by the appearance of white smoke when a stick moistened with hydrochloric acid is introduced. Ammonia and its salts are widely used in industry and technology, agriculture, and everyday life. Their main applications are shown in Figure 31.

Rice. 31. Application of ammonia and ammonium salts:

1-5 - production of mineral fertilizers; 6 - production of nitric acid; 7 - production of explosives; 8 - for soldering; 9 - in refrigeration units; 10 - in medicine and everyday life (ammonia)

6. acids, ionic equation

You have already become acquainted with one of the representatives of substances of this class when you considered volatile hydrogen compounds using the example of hydrogen chloride HCl. Its solution in water is hydrochloric acid. They have the same formula HCl. Similarly, when another volatile substance is dissolved in water hydrogen connection- hydrogen sulfide H2S forms a solution of weak hydrosulfide acid with the formula H2S.

The molecules of these acids consist of two elements, that is, they are binary compounds. However, the class of acids also includes compounds consisting of a larger number of chemical elements. Kick, the third element included in the composition of acid is oxygen. Therefore, such acids are called oxygen-containing, in contrast to HCl and H2S, which are called oxygen-free. Let's list some oxygen-containing acids.

Please note that all acids (oxygen-containing and oxygen-free) necessarily contain hydrogen, which is written in the first place in the formula. The rest of the formula is called the acid residue. For example, in HCl the acidic residue is Cl-.

Acids are complex substances whose molecules consist of hydrogen atoms and acid residues.
As a rule, acidic residues form nonmetallic elements.

Using the formulas of acids, you can determine the oxidation states of the atoms of the chemical elements that form acids.
For binary acids this is easy to do. Since hydrogen has an oxidation state of +1. then in the compound H+1Cl-1, chlorine has an oxidation state of -1, and in the compound H2+1S-2, sulfur has an oxidation state of -2.

It will not be difficult to calculate the oxidation states of the atoms of nonmetal elements that form the acidic residues of oxygen-containing acids. You just need to remember that the total oxidation state of the atoms of all elements in the compound is zero, and the oxidation state of hydrogen is +1 and oxygen is -2.
Knowing the oxidation state of a nonmetal element that forms the acidic residue of an oxygen-containing acid, it is possible to determine which oxide corresponds to it. For example, sulfuric acid HgSO, in which sulfur has an oxidation state of +6, corresponds to sulfur oxide (VI) S03; nitric acid HN03, in which the oxidation state of nitrogen is +5, corresponds to nitrogen oxyl (V) NzOu.

Using the formulas of acids, you can also determine the total charge that acid residues have. The charge of an acid residue is always negative and equal to the number of hydrogen atoms in the acid. The number of hydrogen atoms in an acid is called basicity. For monobasic acids containing one hydrogen atom, for example HCl and HN08, the charges of the acid residues are 1-. For dibasic acids, such as H2SO4 and H2S, the charges of the acid residues are 2-, that is
SO4 2- and S 2-.

There are many acids found in nature: lemon acid in lemons, malic acid in apples, oxalic acid in sorrel leaves. Ants protect themselves from enemies by spraying caustic droplets of formic acid. It is also found in bee venom and stinging nettle hairs.

When grape juice sours, acetic acid is obtained, and when milk sours, lactic acid is obtained. The same lactic acid is formed during sauerkraut and during ensiling of livestock feed. We are well aware of citric and acetic acids, which are often used in everyday life. Vinegar used in food is a solution acetic acid.Many acids are needed in national economy in huge quantities, the production of these substances is called multi-tonnage. These include sulfuric and hydrochloric acids.

Sulfuric acid S2SO4 - a colorless liquid, viscous, like oil, odorless, almost twice as heavy as water. Sulfuric acid absorbs moisture from air and other gases. This property of sulfuric acid is used to dry some gases.

When sulfuric acid is mixed with water, it produces a large number of warmth. If water is poured into sulfuric acid, then the water, not having time to mix with the acid, can boil and splash splashes of sulfuric acid on the face and hands of the worker. To prevent this from happening, when dissolving sulfuric acid, you need to pour it into water in a thin stream and mix.

Sulfuric acid chars wood, leather, and fabrics. If a splinter is placed in a test tube with sulfuric acid, then chemical reaction- the splinter is charred. Now it is clear how dangerous it is for splashes of sulfuric acid to come into contact with human skin and clothing.

Solutions of all acids are acidic, but not a single chemist dares to recognize concentrated acids by taste - this is dangerous. There are more effective and safer ways to detect acids. They, like alkalis, are recognized using indicators.

Add a few drops of litmus solution to the acid solutions. purple. The litmus will turn red. Methyl orange changes when exposed to acids Orange color to red-pink.

But silicic acid, since it is insoluble in water, cannot be recognized this way.

Under normal conditions, acids can be solid (phosphoric H3P04, silicon H2SiO2) and liquid (in pure form the liquid will be sulfuric acid H2SO4).

Gases such as hydrogen chloride HCl, hydrogen bromide HBr, hydrogen sulfide H2S form the corresponding acids in aqueous solutions.

You already know that carbonic H2CO3 and sulfurous H2SO3 acids exist only in aqueous solutions, since they are weak and unstable. They easily decompose into carbon (IV) and sulfur (IV) oxides - CO2 and SO2, respectively, and water. Therefore, it is impossible to isolate these acids in their pure form. The concepts of volatility and stability (stability) are often confused. Volatile acids are acids whose molecules easily pass into a gaseous state, that is, evaporate. For example, hydrochloric acid is a volatile, but stable, stable acid. It is impossible to judge the volatility of unstable acids. For example, non-volatile insoluble silicic acid, when standing, decomposes into water u SiO2. Aqueous solutions of hydrochloric, nitric, sulfuric, phosphoric and a number of other acids are colorless. Aqueous solutions of chromic acid H2CrJ2 are yellow in color, and manganese acid HMnO4 is crimson. However, no matter how different the acids are, they all form hydrogen cations upon dissociation, which determine the series general properties: sour taste, change in color of indicators (litmus and methyl orange), interaction with other substances. The division of acids into groups according to various characteristics is presented in Table 10.

“Structure of the atom and atomic nucleus” - Protons and neutrons. Examples of electronic formulas of atoms. Image of electron orbitals. Calculation of the number of protons, neutrons and electrons. Atom and nucleus. Levels, sublevels and orbitals. Choose the correct answer. Goals. Finding an electron in an atom. Control materials. Write the electronic formula. Opening the core.

“Structure of the nucleus of an atom” - The atom is neutral, because. the charge of the nucleus is equal to the total charge of the electrons. Electrons move around the nucleus. 1919 Rutherford studied the interaction of particles with the nuclei of nitrogen atoms. Foil made of the metal being studied. CONTENTS Module 1 1. Atomic structure. The total number of nucleons in a nucleus is called the mass number and is denoted A.

“Composition of the nucleus of an atom” - The nucleus of an atom of a chemical element. Graph of specific nucleon bonding in a nucleus. Core charge. Scheme of Rutherford's experiments. Dimensions of atomic nuclei. Discovery of the proton. The number of neutrons in the nucleus of an atom. Nuclear forces. Properties of nuclear forces. Proton and neutron. Mass defect. Formula for finding binding energy. Density of nuclear matter.

““Nuclear structure” physics” - How many nucleons do nuclei contain. Helium nucleus. Charge number. New element. The structure of the atomic nucleus. Learn about the history of the discovery of the neutron. Isotopes. A particle that has no charge. Proton-neutron model of the atomic nucleus. Determine the nucleon composition of the nucleus. Neutron. A hero with short arms.

“Composition of the atomic nucleus” - Lesson plan. NUCLEAR FORCES – attractive forces that bind protons and neutrons in the nucleus. Short-range (r = 2.2 * 10-15 m). PROPERTIES They are only forces of attraction. The charge number is equal to the charge of the nucleus, expressed in elementary electric charges. Does not depend on the presence of charge. Nuclear forces. Mass number.

“Structure of the atomic nucleus” - M - mass number - mass of the nucleus, number of nucleons, number of neutrons M-Z. Radioactivity is proof of the complex structure of atoms. Nuclear chain reaction. Was Prometheus right when he gave people fire? Fission of the atomic nucleus. Geiger counter Wilson chamber. The structure of the atomic nucleus. Repeating and generalizing lesson

NITROGEN
N (nitrogenium),
chemical element(at. number 7) VA subgroups 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 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 (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. 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.
See also ATOMIC STRUCTURE.
Molecular nitrogen. From determinations of gas density it has been established that the nitrogen molecule is diatomic, i.e. molecular formula nitrogen has the form NєN (or N2). For two nitrogen atoms, the three outer 2p electrons of each atom form a triple bond:N:::N:, forming electron pairs. Measured interatomic distance N-N equals 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 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 - 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.

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 quite 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. Connections with such a structure interact with electrical 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 by 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 with the formation of 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 its low concentration is 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 .

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 foreign words 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 system, 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