Meet nitrogen and its cycle

The proposed series of articles examines the logical chemistry of nitrogen - the prediction of the physical and chemical properties of substances based on the structure of atoms and molecules. Additional information in numbers is provided that goes beyond the scope of the school course. Environmental chemistry of nitrogen contains information about how living and nonliving nature reacts to specific substances. Recommendations are given to reduce the harmful effects of these substances.

"Life Denier"

The history of the discovery of nitrogen is closely intertwined with the history of the development of chemistry in general. Since the 17th century Scientists began to become interested in gases, and instruments were invented to collect and study them. Naturally, they became interested in air - a gas that is always at hand. Karl Scheele, Joseph Priestley, Daniel Rutherford, and Henry Cavendish were studying air at about the same time. Soon they released nitrogen from the air. As a matter of fact, the nitrogen was released on its own.

For example, an analysis of the gas mixture formed when a candle is burned under a hood (Fig. 1) shows the following. Oxygen from the air is consumed during combustion, carbon dioxide reacts with lime, and water vapor is absorbed by calcined calcium chloride. What remains is nitrogen (and a “bouquet” of inert gases in an amount of about one percent by volume). It was not difficult to isolate nitrogen gas even at that time. Discovering a new simple substance was much more difficult.

Rice. 1. Perhaps this is how nitrogen was obtained and its properties were tested

The processes of interaction of substances with air were explained by the theory of phlogiston. It was assumed that phlogiston is a flammable substance that is released from substances during combustion and disperses into the air. This theory is very similar to modern theories of combustion, exactly the opposite, as if upside down. It was phlogiston that prevented nitrogen from being recognized as a simple substance. Antoine Laurent Lavoisier turned the theory upside down and freed science from phlogiston. Along the way, he convinced scientists that nitrogen is a simple substance, and not a compound of something with phlogiston.

In an atmosphere of pure nitrogen, animals and plants die. Lavoisier knew about this and called nitrogen “denying life” (from the Greek a - negative particle and Zoos - living). Element No. 7 has this name in French and Russian. The German name has a similar meaning: Stickstoff – asphyxiating substance. The British took a different path, adopting the name Nitrogen - saltpeter that gives birth.

The substance, without which life is impossible, was called “life-denying” (without nitrogen, life is impossible, if only because it dilutes the oxygen in the air, because life is just as impossible in pure oxygen as in pure nitrogen). Even the name showed the uniqueness of element No. 7.

Nitrogen cycle

Nitrogen is one of the elements actively used by nature to build substances. The element nitrogen is found in you and me and in almost all objects of fauna and flora. The role of fauna and flora in the nitrogen cycle is different. Fauna eats and transforms those substances that flora has made from the “bricks” of inanimate nature. The “bricks” themselves are air and aqueous solutions of various substances. This “division of labor” is a very wise creation of nature. It allows us to create more complex and dynamic organisms.

The nitrogen cycle in animals is the domain of biochemistry, and we'll look at the inorganic chemistry in and around a green blade of grass.

Nitrogen enters the plant in the form of ions dissolved in water. These can be ammonium ions, which are immediately used for their intended purpose, or nitrate ions, which must be converted into ammonium ions in order to be absorbed by the body. The main direction of rearrangement of nitrogen-containing compounds in a plant: nitrates ® nitrites ® ammonium ions ® amines ® amino acids ® proteins (Fig. 2). Nature provides the possibility of a reverse motion if any of the substances is in excess. This will be discussed in more detail when discussing the environmental properties of nitrogen compounds.

Rice. 2. The nitrogen cycle in and around a green blade of grass (the sign shows organic compounds - amines and proteins)

In the soil, with the help of bacteria, proteins that enter it are converted through amino acids and amines into nitrate ions as the least toxic and most stable. Atmospheric nitrogen is also involved in this circle. Some bacteria convert N2 molecules into ammonium ions, while other bacteria return N2 nitrogen to the atmosphere, reducing it from nitrate ions. Some of the atmospheric nitrogen is converted into compounds by lightning during thunderstorms. Nature, unlike humans, uses atmospheric nitrogen very economically and rationally.

Man actively interferes with the cycle of all substances and elements, including nitrogen. Now the balance has changed: the simple substance nitrogen – nitrogen compounds. Many nitrogenous compounds, instead of performing useful functions, are spent only on poisoning the environment.

Mainly, these are compounds contained in emissions from various enterprises and formed during fuel combustion.

In addition to the fact that man throws many compounds down the drain, he also separated the two halves of the nitrogen cycle. The “nitrate ® protein” chain is carried out in the fields, and the “protein ® nitrate” chain is carried out in cities and towns. Therefore, there was a need to add nitrogen to the soil, fertilize it, make it “good” and fertile (Fig. 3). This is not an easy task. You can't just dump a bag of ammonium nitrate on the garden bed. First you need to carefully find out what exactly a particular soil needs. It is necessary to love and know your land. And it doesn’t matter whether it’s endless fields or a flowerbed under the windows.

Rice. 3. Divided nitrogen cycle:

– nitrogen compounds emitted into the atmosphere by car engines, factories, etc. (oxygen-containing and oxygen-free compounds, ammonia, cyanogen, etc.);
– nitrogen compounds separated by distance, included in the cycle only with the help of technology;
– ordinary vehicles transporting nitrogen-containing compounds

Of course, the idea of ​​the main compounds in the “circle of transformations” of nitrogen (see Fig. 2) turned out to be very schematic. In addition to ammonium and nitrate ions, the plant absorbs and processes other nitrogen compounds. The transformations of some proteins into others are interesting and complex. But it is impossible to embrace the immensity.

Features of the structure of the nitrogen atom

The address of nitrogen in the periodic table of chemical elements is the second period, fifth group, main subgroup. The basis for the uniqueness of the element nitrogen lies in the structure of the second electron layer. This layer can accommodate a maximum of eight electrons in four atomic orbitals (one s and three p).

The difference between the second layer and the third and more distant electronic layers from the nucleus of the atom is that it does not have spare d-orbitals. Therefore, the maximum number of chemical bonds for elements of the second period is four. (For an electron to move to the third layer, much more energy is required than can be released during the formation of a chemical bond. This is why pentavalent nitrogen does not exist in nature.)

Elements of the second period have one more feature: their outermost layer is occupied by only two electrons. This means that when two atoms come closer together to form single and double bonds, their electron shells will repel each other less.

We looked at how the presence of an element in the 2nd period of the periodic table affects the structure of its atom. Now let’s look at what staying in Group V entails. Elements of group Va have five electrons in the outer electron layer. An octet (up to eight) is missing three electrons. It can be assumed that the simple substance nitrogen has a molecular structure. Indeed, N 2 (NєN) molecules consist of two atoms connected by a triple bond.

Substances with a molecular crystal lattice have relatively low melting and boiling points, and the forces of intermolecular attraction are several orders of magnitude weaker than chemical covalent bonds. On top of that, the nitrogen molecule is very light. The conclusion suggests itself that the melting and boiling points of nitrogen are not even relatively, but significantly low. (Indeed, the air in Antarctica does not liquefy on its own, although temperatures there can reach -80 °C.)

The nitrogen molecule is formed by identical atoms and is therefore non-polar. Therefore, nitrogen is slightly soluble in water. However, one should remember that the solubility of nitrogen in water increases with increasing external pressure. That is why divers from great depths, where the pressure is higher, have to rise slowly. Otherwise, the nitrogen dissolved in the blood, released, boils, as it were, forming bubbles in the blood vessels.

Knowledge about the structure of the nitrogen molecule can also help in predicting the chemical properties of this substance. During chemical reactions, existing bonds between atoms are broken and new ones are formed. It is clear that breaking a triple bond is much more difficult than a double and single bond (a twisted rope is more difficult to break than one of the threads). Most likely, the simple substance nitrogen should be reluctant to interact with other substances.

This is how it really is. Under normal conditions and even with slight heating, nitrogen practically does not react with anything. And that's great, because otherwise we wouldn't be on this blue planet and the planet might not be blue.

Questions. Make suggestions about life on Earth if nitrogen were a solid, like carbon, phosphorus, and silicon. What happens if nitrogen becomes as active as oxygen and fluorine?

Physical properties of nitrogen in numbers

Molar mass – 28 g/mol. The melting point is –210 °C, the boiling point is –195.8 °C. The density of nitrogen gas under normal conditions (1 atm, 0 °C) is 1.251 g/l. (For reference: the density of air under normal conditions is 1.293 g/l, nitrogen is slightly lighter than air.) The density of liquid nitrogen (at –196 °C) is 0.808 g/cm3. Solubility at 1 atm and 0 °C - 2.35 ml of gas per 100 g of water, at 20 °C - 1.54 ml of N 2 per 100 g of water.

Chemical properties of nitrogen

1. Reactions with metals.

Under normal conditions, nitrogen reacts with lithium:

When heated, reactions occur with Na, Ca, Mg, Mn. Manganese, for example, combines with nitrogen at 1200 °C:

3Mn + N 2 = Mn 3 N 2.

In other words, nitrogen reacts only with the most active metals, and even then only reluctantly.

2. Reactions with non-metals.

When heated to 1200 °C, nitrogen begins to react with oxygen. However, under these conditions little nitric oxide is produced. At 3000 °C, equilibrium in the reaction is established instantly and nitrogen oxide is formed in noticeable quantities:

This temperature is possible in the lightning channel, so it is during thunderstorms that plants naturally replenish their supply of nitrogen compounds.

When heated under pressure (500 °C, 300 atm) in the presence of a catalyst (for example, iron activated by calcium and aluminum oxides), nitrogen reacts with hydrogen. Even under such harsh conditions, the ammonia yield does not exceed 30%, but this is sufficient for the industrial use of this reaction:

3. Reactions with complex substances.

When calcium carbide is heated to 1000 °C in a tightly closed furnace with nitrogen supplied there under pressure, the following reaction occurs between them:

Ecological properties of the simple substance nitrogen

Nitrogen, a “life-denying element,” is actually a life-affirming element. And first of all, nitrogen affirms life with its inertness. By diluting oxygen, it allows the oxidation reactions of organic substances to proceed relatively slowly and stop at “half stations” - intermediate degrees of carbon oxidation. For all living things, nitrogen is harmless in any quantity (provided, however, that the necessary quantities of oxygen and carbon dioxide are present).

At the same time, nitrogen can be a carrier of harmful and simply dangerous. For example, nitrogen can easily blow off the roof of a house during a hurricane. After all, air masses - a hurricane, a monsoon, and just wind - are all three-quarters nitrogen. And the ecological chemistry of nitrogen turns into the ecology of air - a very broad topic. Therefore, it’s time to put an end to it.

DEFINITION

Nitrogen- the seventh element of the Periodic Table. Designation - N from the Latin "nitrogenium". Located in the second period, VA group. Refers to non-metals. The nuclear charge is 7.

Most of the nitrogen is in a free state. Free nitrogen is the main component of air, which contains 78.2% (vol.) nitrogen. Inorganic nitrogen compounds do not occur in nature in large quantities, with the exception of sodium nitrate NaNO 3, which forms thick layers on the Pacific coast of Chile. The soil contains small amounts of nitrogen, mainly in the form of nitric acid salts. But in the form of complex organic compounds - proteins - nitrogen is part of all living organisms.

In the form of a simple substance, nitrogen is a colorless, odorless gas and very slightly soluble in water. It is slightly lighter than air: the mass of 1 liter of nitrogen is 1.25 g.

Atomic and molecular mass of nitrogen

The relative atomic mass of an element is the ratio of the mass of an atom of a given element to 1/12 of the mass of a carbon atom. Relative atomic mass is dimensionless and is denoted by A r (the index “r” is the initial letter of the English word relative, which means “relative”). The relative atomic mass of atomic nitrogen is 14.0064 amu.

The masses of molecules, as well as the masses of atoms, are expressed in atomic mass units. The molecular mass of a substance is the mass of a molecule, expressed in atomic mass units. The relative molecular mass of a substance is the ratio of the mass of a molecule of a given substance to 1/12 of the mass of a carbon atom, the mass of which is 12 amu. It is known that the nitrogen molecule is diatomic - N 2. The relative molecular weight of a nitrogen molecule will be equal to:

M r (N 2) = 14.0064× 2 ≈ 28.

Isotopes of nitrogen

In nature, nitrogen exists in the form of two stable isotopes 14 N (99.635%) and 15 N (0.365%). Their mass numbers are 14 and 15, respectively. The nucleus of an atom of the nitrogen isotope 14 N contains seven protons and seven neutrons, and the isotope 15 N contains the same number of protons and six neutrons.

There are fourteen artificial isotopes of nitrogen with mass numbers from 10 to 13 and from 16 to 25, of which the most stable isotope 13 N with a half-life of 10 minutes.

Nitrogen ions

The outer energy level of the nitrogen atom has five electrons, which are valence electrons:

1s 2 2s 2 2p 3 .

The structure of the nitrogen atom is shown below:

As a result of chemical interaction, nitrogen can lose its valence electrons, i.e. be their donor, and turn into positively charged ions or accept electrons from another atom, i.e. be their acceptor and turn into negatively charged ions:

N 0 -5e → N 2+ ;

N 0 -4e → N 4+ ;

N 0 -3e → N 3+ ;

N 0 -2e → N 2+ ;

N 0 -1e → N 1+ ;

N 0 +1e → N 1- ;

N 0 +2e → N 2- ;

N 0 +3e → N 3- .

Nitrogen molecule and atom

The nitrogen molecule consists of two atoms - N 2. Here are some properties characterizing the nitrogen atom and molecule:

Examples of problem solving

EXAMPLE 1

Exercise	To form ammonium chloride, 11.2 liters (n.s.) of ammonia gas and 11.4 liters (n.s.) of hydrogen chloride were taken. What is the mass of the reaction product formed?
Solution	Let us write the equation for the reaction of producing ammonium chloride from ammonia and hydrogen chloride: NH 3 + HCl = NH 4 Cl. Let's find the number of moles of starting substances: n(NH 3) = V(NH 3) / V m; n(NH 3) = 11.2 / 22.4 = 0.5 mol. n(HCl) = V(NH 3) / V m; n(HCl) = 11.4 / 22.4 = 0.51 mol. n(NH 3) n(NH 4 Cl) = n(NH 3) = 0.5 mol. Then, the mass of ammonium chloride will be equal to: M(NH 4 Cl) = 14 + 4 × 1 + 35.5 = 53.5 g/mol. m(NH 4 Cl) = n(NH 4 Cl) × M(NH 4 Cl); m(NH 4 Cl) = 0.5 × 53.5 = 26.75 g.
Answer	26.75 g

EXAMPLE 2

Exercise	10.7 g of ammonium chloride was mixed with 6 g of calcium hydroxide and the mixture was heated. What gas and how much of it was released by mass and volume (n.s.)?
Solution	Let us write the reaction equation for the interaction of ammonium chloride with calcium hydroxide: 2NH 4 Cl + Ca(OH) 2 = CaCl 2 + 2NH 3 - + 2H 2 O. Let us determine which of the two reactants is in excess. To do this, we calculate their number of moles: M(NH 4 Cl) = A r (N) + 4×A r (H) + A r (Cl); M(NH 4 Cl) = 14 + 4×1 + 35.5 = 53.5 g/mol. n(NH 4 Cl) = m (NH 4 Cl) / M(NH 4 Cl); n(NH 4 Cl) = 10.7 / 53.5 = 0.1 mol. M(Ca(OH) 2) = A r (Ca) + 2×A r (H) + 2×A r (O); M(Ca(OH) 2) = 40 + 2×1 + 2×16 = 42 + 32 = 74 g/mol. n(Ca(OH) 2) = m (Ca(OH) 2) / M(Ca(OH) 2); n(Ca(OH) 2) = 6 / 74 = 0.08 mol. n(Ca(OH)2) n(NH 3) = 2×n(Ca(OH) 2) = 2×0.08 = 0.16 mol. Then, the mass of ammonia will be equal to: M(NH 3) = A r (N) + 3×A r (H) = 14 + 3×1 = 17 g/mol. m(NH 3) = n(NH 3) × M(NH 3) = 0.16 × 17 = 2.72 g. The volume of ammonia is equal to: V(NH 3) = n(NH 3) ×V m; V(NH 3) = 0.16 × 22.4 = 3.584 l.
Answer	As a result of the reaction, ammonia was formed with a volume of 3.584 liters and a mass of 2.72 g.

Nitrogen is a chemical element that is known to everyone. It is designated by the letter N. It can be said to be the basis of inorganic chemistry, and therefore it begins to be studied in the eighth grade. In this article we will take a closer look at nitrogen, as well as its characteristics and properties.

History of element discovery

Compounds such as ammonia, nitrate, and nitric acid were known and used in practice long before pure nitrogen was obtained in a free state.

In an experiment conducted in 1772, Daniel Rutherford burned phosphorus and other substances in a glass bell. He found that the gas remaining after the combustion of compounds does not support combustion and respiration, and called it “suffocating air.”

In 1787, Antoine Lavoisier established that the gases that make up ordinary air are simple chemical elements, and proposed the name “Nitrogen”. A little later (in 1784), physicist Henry Cavendish proved that this substance is part of nitrate (a group of nitrates). This is where the Latin name for nitrogen comes from (from the Late Latin nitrum and Greek gennao), proposed by J. A. Chaptal in 1790.

By the beginning of the 19th century, scientists had clarified the chemical inertness of the element in a free state and its exceptional role in compounds with other substances. From that moment on, the “binding” of air nitrogen became the most important technical problem in chemistry.

Physical properties

Nitrogen is slightly lighter than air. Its density is 1.2506 kg/m³ (0 °C, 760 mm Hg), melting point - -209.86 °C, boiling point - -195.8 °C. Nitrogen is difficult to liquefy. Its critical temperature is relatively low (-147.1 °C), while the critical pressure is quite high - 3.39 Mn/m². Density in liquid state - 808 kg/m³. This element is less soluble in water than oxygen: 23.3 g of N can be dissolved in 1 m³ (at 0 °C) of H₂O. This figure is higher when working with some hydrocarbons.

When heated to low temperatures, this element interacts only with active metals. For example, with lithium, calcium, magnesium. Nitrogen reacts with most other substances in the presence of catalysts and/or at high temperatures.

The compounds of N with O₂ (oxygen) N₂O₅, NO, N₂O₃, N₂O, NO₂ have been well studied. From them, during the interaction of elements (t - 4000 ° C), NO oxide is formed. Further, during the cooling process, it is oxidized to NO₂. Nitrogen oxides are formed in the air during the passage of atmospheric discharges. They can be obtained by the action of ionizing radiation on a mixture of N and O₂.

When N₂O₃ and N₂O₅ are dissolved in water, respectively, the acids HNO₂ and HNO₂ are obtained, forming salts - nitrates and nitrites. Nitrogen combines with hydrogen exclusively in the presence of catalysts and at high temperatures, forming NH₃ (ammonia). In addition, other (they are quite numerous) compounds of N with H₂ are known, for example diimide HN = NH, hydrazine H₂N-NH₂, octazone N₈H₁₄, acid HN₃ and others.

It is worth saying that most hydrogen + nitrogen compounds are isolated exclusively in the form of organic derivatives. This element does not react (directly) with halogens, so all its halides are obtained only indirectly. For example, NF₃ is formed when ammonia reacts with fluorine.

Most nitrogen halides are weakly stable compounds; oxyhalides are more stable: NOBr, NO₂F, NOF, NOCl, NO₂Cl. Direct combination of N with sulfur also does not occur; N₄S₄ is obtained during the reaction of ammonia + liquid sulfur. When hot coke reacts with N, cyanogen (CN)₂ is formed. By heating acetylene C₂H₂ with nitrogen to 1500 °C, hydrogen cyanide HCN can be obtained. When N interacts with metals at relatively high temperatures, nitrides are formed (for example, Mg₃N₂).

When ordinary nitrogen is exposed to electric discharges [at a pressure of 130–270 n/m² (corresponding to 1–2 mm Hg)] and during the decomposition of Mg₃N₂, BN, TiNx and Ca₃N₂, as well as during electric discharges in the air, active nitrogen can be formed, having increased energy reserves. It, unlike the molecular one, interacts very energetically with hydrogen, sulfur vapor, oxygen, some metals and phosphorus.

Nitrogen is part of quite a few important organic compounds, including amino acids, amines, nitro compounds and others.

Getting nitrogen

In the laboratory, this element can be easily obtained by heating a concentrated solution of ammonium nitrite (formula: NH₄NO₂ = N₂ + 2H₂O). The technical method for obtaining N is based on the separation of pre-liquefied air, which is subsequently subjected to distillation.

Application area

The main part of the free nitrogen obtained is used in the industrial production of ammonia, which is then processed in fairly large quantities into fertilizers, explosives, etc.

In addition to the direct synthesis of NH₃ from elements, the cyanamide method developed at the beginning of the last century is used. It is based on the fact that at t = 1000 °C calcium carbide (formed by heating a mixture of coal and lime in an electric furnace) reacts with free nitrogen (formula: CaC₂ + N₂ = CaCN₂ + C). The resulting calcium cyanamide decomposes under the influence of heated water vapor into CaCO₃ and 2NH₃.

In its free form, this element is used in many industries: as an inert medium in various metallurgical and chemical processes, when pumping flammable liquids, for filling space in mercury thermometers, etc. In its liquid state, it is used in various refrigeration units. It is transported and stored in steel Dewar vessels, and compressed gas is stored in cylinders.

Many nitrogen compounds are also widely used. Their production began to develop rapidly after the First World War and has now reached truly enormous proportions.

This substance is one of the main biogenic elements and is part of the most important elements of living cells - nucleic acids and proteins. However, the amount of nitrogen in living organisms is small (approximately 1–3% by dry weight). The molecular material present in the atmosphere is assimilated only by blue-green algae and some microorganisms.

Quite large reserves of this substance are concentrated in the soil in the form of various mineral (nitrates, ammonium salts) and organic compounds (composed of nucleic acids, proteins and their breakdown products, including not yet completely decomposed remains of flora and fauna).

Plants perfectly absorb nitrogen from the soil in the form of organic and inorganic compounds. Under natural conditions, special soil microorganisms (ammonifiers) are of great importance, which are capable of mineralizing soil organic N to ammonium salts.

Nitrate nitrogen in the soil is formed during the life of nitrifying bacteria, discovered by S. Winogradsky in 1890. They oxidize ammonium salts and ammonia to nitrates. Part of the substance assimilated by flora and fauna is lost due to the action of denitrifying bacteria.

Microorganisms and plants perfectly absorb both nitrate and ammonium N. They actively convert inorganic material into various organic compounds - amino acids and amides (glutamine and asparagine). The latter are part of many proteins of microorganisms, plants and animals. The synthesis of asparagine and glutamine by amidation (enzymatic) of aspartic and glutamic acids is carried out by many representatives of flora and fauna.

The production of amino acids occurs through the reductive amination of a number of keto acids and aldehyde acids, resulting from enzymatic transamination, as well as from the oxidation of various carbohydrates. The end products of ammonia (NH₃) assimilation by plants and microorganisms are proteins, which are part of the cell nucleus, protoplasm, and are also deposited in the form of so-called storage proteins.

Humans and most animals can synthesize amino acids only to a fairly limited extent. They are not able to produce eight essential compounds (lysine, valine, phenylalanine, tryptophan, isoleucine, leucine, methionine, threonine), and therefore their main source of nitrogen is proteins consumed with food, that is, ultimately, the own proteins of microorganisms and plants.

MOBUSOSH No. 2

Abstract in chemistry on the topic:

“Characteristics of elements of the nitrogen subgroup”

Prepared by: Nasertdinov K.

Checked:

Agidel-2008

1. Characteristics of elements of the nitrogen subgroup

2. Structure and characteristics of atoms

2.1 Nitrogen

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

2.3.1 Properties of nitric acid

2.3.2 Salts of nitric acid and their properties

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 of the 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, and in the mining 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 by the 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
	N 2	Molecular	Gas without color, smell, taste, soluble in water
Phosphorus white	P 4	Tetrahedral molecule. Molecular crystal lattice.	Soft solid, colorless, slightly soluble in water, soluble in carbon sulfur
Arsenic gray	As 4		Brittle crystalline substance with metal. shine on a fresh break. Insoluble in water. Very weak conductor of electricity		Sublimates, passes from solid to gaseous (steam) at 615 o C
	Sb 4		Silvery-white crystalline substance, brittle, poor conductor of heat and electricity
	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

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 valence. The electronic structure is represented by seven electrons located at two energy levels. Electronic formula: 1s 2 2s 2 2p 3. Nitrogen oxidation states: - 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

Under normal conditions, nitrogen is a 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 for the production of nitric acid and nitrogen-containing mineral fertilizers, salts, and 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.

NO 2	Nitric oxide (IV) - nitrogen dioxide	Salt-forming	Brown gas with a specific odor, soluble in water, easily dimerizes
N2O5	Nitric oxide (V) – nitric anhydride		White crystal- lic substance. t pl = 32.3 o C, soluble in water.		Shows properties of acid oxides, thermally unstable, toxic

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 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, highly soluble in water, and 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.

Qualitative reaction to solid nitrates: a pinch of salt is thrown into the fire of the burner, and if the salt is a nitrate, then a bright flash will occur due to the decomposition of the salt with the release of oxygen.

2.3.3 Use of nitric acid and its salts

Nitric acid is one of the large-scale and important products of the chemical industry. It is widely used for the production of fertilizers, smokeless powder, explosives, medicines, dyes, and plastics. Its salts are used in pyrotechnics; for the production of fertilizers, explosives, and some metal oxides.

2.4 Phosphorus

Phosphorus is a non-metal element. In terms of the number of electrons and electronic configuration (3s 2 3p 3), the phosphorus atom is an analogue of nitrogen. But compared to the nitrogen atom, the phosphorus atom has a larger radius, lower ionization energy and OEO, so phosphorus exhibits weaker non-metallic properties. Oxidation states: - 3,+3,+5.

Phosphorus in the free state forms allotropic modifications: white, red and black phosphorus. Allotropic modifications are interrelated and can transform into each other. Phosphorus in reactions can be both a reducing agent and an oxidizing agent. In reactions with active metals, phosphorus acquires an oxidation state of 3.

The reaction products are phosphides (unstable compounds, easily decomposed by water to form PH 3.

Allotropic forms	Composition designation	Lattice type	Characteristics of the most important properties
White phosphorus	P 4	Molecular lattice	A white crystalline substance with a yellowish tint and a garlicky odor; t pl =44 o C, t boil =280 o C, t pl =40 o C (in crushed form). Well soluble in carbon disulfide. Glows in the dark. Poisonous!
Red phosphorus		Atomic lattice	Red-brown powder, odorless, insoluble in water and carbon disulfide; t pl. =260 o C, t pl. does not, because before melting it turns into white phosphorus vapor. Doesn't light up. Non-toxic, non-volatile.
Black phosphorus		Atomic lattice	A substance similar to graphite. Black, greasy to the touch, heavier than white and red phosphorus; t flammable >490 o C. Insoluble in water and carbon sulfur. Semiconductor. Doesn't light up. Non-toxic, non-volatile

Table of allotropic forms of phosphorus.

2.4.1 Phosphorus compounds

The compound of phosphorus with hydrogen is gaseous hydrogen phosphide, or phosphine PH 3 (a colorless, poisonous gas with a garlicky odor, flammable in air).

Phosphorus has several oxides: phosphorus oxide (III) P 2 O 3 (a white crystalline substance, formed during the slow oxidation of phosphorus in conditions of lack of oxygen, toxic) and phosphorus oxide (V) P 2 O 5 (formed from P 2 O 3 when it heating, soluble in water to form phosphorous acid of medium strength) the most important. The most characteristic property of the second is hygroscopicity (absorption of water vapor from the air), while it spreads into an amorphous mass of HPO 3. When P 2 O 5 is boiled, phosphoric acid H 3 PO 4 is formed (white crystalline substance, melts in air, t pl = 42.35 o C, non-toxic, soluble in water, electrolyte, obtained by oxidizing 32% nitric acid) . Phosphates of almost all metals (except alkali) are insoluble in water. Dihydrogen phosphates are highly soluble in water.

2.4.2 Application of phosphorus and its compounds

A large amount of phosphorus is used in the production of matches, white phosphorus is widely used in the creation of incendiary shells, smoke bombs, shells and bombs, salts of phosphoric acid are used in agriculture as phosphorus fertilizers.

2.5 Mineral fertilizers

Type and name	Chem. compound	Condition and appearance	Nutrient element and its content, %
Nitrogen fertilizers
Sodium nitrate (Chilean saltpeter)	NaNo 3	White-gray crystalline substance, hygroscopic, soluble in H 2 O
Ammonium nitrate	NH4NO3	White crystalline, very hygroscopic substance
Ammonium sulfate	(NH 4 ) 2 SO 4	White-gray crystalline powder, slightly hygroscopic
Urea (urea)	(NH 2 ) 2 CO	White fine-crystalline hygroscopic substance
Liquid concentrated ammonia	NH 3	Liquid with a pungent odor, highly soluble in water
Ammonia water	NH 3 +H 2 O	Ammonia solution in water
Ammonia	NH 4 NO 3 + NH 3 + H 2 O	Aqueous solution of ammonium nitrate and ammonia
Phosphorus fertilizers			P2O5
Simple superphosphate	Ca (H 2 PO 4 ) 2 x x CaSO 4	Gray powdery substance, soluble in water with CaSO 4 ballast
Double superphosphate	Ca(H2PO4)2	Similar to simple superphosphate, but without ballast.
Precipitate	CaHPO 4 x x 2H 2 O	White-gray powdery substance, highly soluble in water
Potash fertilizers			K2O
Potassium chloride		White, finely crystalline substance, highly soluble in water
Potassium sulfate	K2SO4	White crystalline non-hygroscopic substance
Complex fertilizers
Potassium nitrate	KNO 3	White crystalline substance highly soluble in water	Double fertilizer K and N
	NH4H2PO4		P 2 O 5 -46-50%
Diammofos	(NH 4 ) 2 HPO 4		N-21%, P 2 O 5 -53%
Ammofoska	(NH 4 ) 2 HPO 4 + NH nitrogen in nature Total characteristic elements subgroups nitrogen PROPERTIES NITROGEN Isotopes, ... oxygen in water. General characteristic elements subgroups nitrogen Nitrogen phosphorus Arsenic Antimony Bismuth Structure... Metals by-product subgroups Group I Coursework >> Chemistry Heated copper is exposed to oxides nitrogen: N2O and NO interact..., and various physico-chemical characteristics. Silver compounds have a significant... in this regard, somewhat closer than others elements subgroups copper is related to alkali metals... Periodic table elements Mendeleev Abstract >> Chemistry ... elements can vary over a wide range from non-metallic to metallic (for example, in the main subgroup Group V nitrogen... and etc.). So, the main thing characteristic an atom is not the atomic mass... and the exact characteristic atom, and therefore element. From... Chemical elements, their connections and valency Test >> Chemistry Of the most important characteristics element. More than 110 chemicals are known Elements, they, ... electronegativity. There are chemical elements main subgroups, or intransitive elements, in which... in numbers: Nitrous nitrogen N2O Oxide nitrogen NO Nitrous anhydride... Chemical element- Scandium Abstract >> Chemistry Fe3+, Mn3+), elements subgroups Al, Be, and also elements yttrium subgroups, together with which... (450 °C) the hydride ScH2 is formed, with nitrogen(600-800 °C) - ScN nitride, ... with beryllium, having unique characteristics in strength and heat resistance. So...

Nitrogen is a chemical element with atomic number 7. It is an odorless, tasteless and colorless gas.


Thus, a person does not feel the presence of nitrogen in the earth’s atmosphere, while it consists of 78 percent of this substance. Nitrogen is one of the most common substances on our planet. You can often hear that without nitrogen there would be no food, and this is true. After all, the protein compounds that make up all living things necessarily contain nitrogen.

Nitrogen in nature

Nitrogen is found in the atmosphere in the form of molecules consisting of two atoms. In addition to the atmosphere, nitrogen is found in the Earth's mantle and in the humus layer of the soil. The main source of nitrogen for industrial production is minerals.

However, in recent decades, when mineral reserves began to deplete, an urgent need arose to separate nitrogen from the air on an industrial scale. This problem has now been solved, and huge volumes of nitrogen for industrial needs are extracted from the atmosphere.

The role of nitrogen in biology, the nitrogen cycle

On Earth, nitrogen undergoes a number of transformations in which both biotic (life-related) and abiotic factors are involved. Nitrogen enters plants from the atmosphere and soil, not directly, but through microorganisms. Nitrogen-fixing bacteria retain and process nitrogen, converting it into a form that can be easily absorbed by plants. In the plant body, nitrogen is converted into complex compounds, in particular proteins.

Through the food chain, these substances enter the bodies of herbivores and then predators. After the death of all living things, nitrogen returns to the soil, where it undergoes decomposition (ammonification and denitrification). Nitrogen is fixed in the soil, minerals, water, enters the atmosphere, and the circle repeats.

Application of nitrogen

After the discovery of nitrogen (this happened in the 18th century), the properties of the substance itself, its compounds, and the possibility of using it on the farm were well studied. Since the reserves of nitrogen on our planet are huge, this element has become extremely actively used.


Pure nitrogen is used in liquid or gaseous form. Liquid nitrogen has a temperature of minus 196 degrees Celsius and is used in the following areas:

— in medicine. Liquid nitrogen is a refrigerant in cryotherapy procedures, that is, cold treatment. Flash freezing is used to remove various tumors. Tissue samples and living cells (in particular, sperm and eggs) are stored in liquid nitrogen. Low temperature allows the biomaterial to be preserved for a long time, and then thawed and used.

The possibility of storing entire living organisms in liquid nitrogen, and, if necessary, defrosting them without any harm, was expressed by science fiction writers. However, in reality it has not yet been possible to master this technology;

— in the food industry Liquid nitrogen is used when bottling liquids to create an inert environment in the container.

In general, nitrogen is used in areas where a gaseous environment without oxygen is required, e.g.

— in fire fighting. Nitrogen displaces oxygen, without which combustion processes are not supported and the fire goes out.

Nitrogen gas has found application in the following industries:

— food production. Nitrogen is used as an inert gas medium to maintain the freshness of packaged products;

— in the oil industry and mining. Pipelines and tanks are purged with nitrogen, it is injected into mines to form an explosion-proof gas environment;

— in aircraft manufacturing The chassis tires are inflated with nitrogen.

All of the above applies to the use of pure nitrogen, but do not forget that this element is the starting material for the production of a mass of various compounds:

- ammonia. An extremely sought-after substance containing nitrogen. Ammonia is used in the production of fertilizers, polymers, soda, and nitric acid. It is itself used in medicine, in the manufacture of refrigeration equipment;

— nitrogen fertilizers;

- explosives;

- dyes, etc.


Nitrogen is not only one of the most common chemical elements, but also a very necessary component used in many branches of human activity.