The highest oxidation state of nitrogen in the compound is equal to. Oxidation states of nitrogen in ammonium

The VA subgroup is formed by p-elements: nitrogenN, phosphorus

P, arsenicAs, antimonySb and bismuthBi.

Elements N, P are typical non-metals,

for nonmetals As and Sb some properties appear

inherent in metals, bismuth has metallic properties

predominate, although it is not a typical metal.

The general formula for valence electrons in elements is

com VA-group –ns 2 np 3.

throne Due to three unpaired electrons all elements in simple substances form three covalent bonds, but in nitrogen three bonds unite 2 atoms, forming a very strong

molecule N N, and for other elements, each atom is connected with three others to form molecules of the E4 type (white

yellow phosphorus and yellow arsenic) or polymer structures.

In nitrogen, a simple substance in any state of aggregation consists of individual molecules , under normal conditions it is a gas. All other elements have simple substances

– hard.

The oxidation state (–3) for elements of the VA group is minimal. It is most stable in N, at

transition to Bi with an increase in the number of electronic layers, its stability increases

gives. The elements N, P, As, Sb with hydrogen form hydrides of the EN3 type,

exhibiting basic properties, they are most pronounced in ammonia-

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ka NH3. In the subgroup, the stability of EN3 compounds and their main properties

va are decreasing.

All elements of the VA group exhibit the highest oxidation state of +5.

All of them form oxides of the E2 O5 type (Bi 2 O 5 oxide is unstable), which correspond to acids the strength of acids weakens when moving down the sub-

The oxidation state +5 is most stable in P . Bi(+5) compounds –

very strong oxidizing agents. Nitric acid, especially concentrated acid, exhibits strong oxidizing properties.

Bismuth has a more stable oxidation state (+3), which is also quite stable in Sb and As. N(+3) compounds, and especially

P(+3), exhibit strong reducing properties.

In the oxidation state +3, all elements of the VA group form oxides

type E 2 O 3. The oxides N and P correspond to weak acids. Oxides and hydroxy-

As and Sb oxides are amphoteric, the basic character predominates in the oxide and hydroxy-

yes Bi(+3). Thus, in the subgroup the acidic nature of oxides and hydro-

oxides of elements in the oxidation state (+3) weakens and increases

basic properties more typical of metal hydroxides.

Elements of group VA, in addition to the listed oxidation states

5, +3, –3, also exhibit other intermediate oxidation states.

For nitrogen, all oxidation states from –1 to +5 are known.

Nitrogen, like all elements of the second period, differs significantly from its electronic analogues . For this reason, and also because of a large number of oxidation states and a variety of compounds, the chemistry of nitrogen is considered

is separated from other elements of the VA-subgroup.

The most common element of the VA group in nature is

phosphorus is present. Its content in the earth's crust is 0.09 mass. %; phosphorus finds-

mainly in the form of calcium phosphate. Nitrogen content – 0.03%, os-

its new share is concentrated in the atmosphere in the form of N2. Nitrogen content in

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air by volume is ~ 78%. In very small quantities in the earth

On the bark, sodium and potassium nitrates (saltpeter) are found. Arsenic, antimony and bismuth are rare elements with a content in the earth's crust of 10–5 5. 10–

4%; in nature they are found mainly in the form of sulfides.

Nitrogen and phosphorus are very important elements of the biosphere, therefore

The main part of the nitrates and phosphorus produced in the chemical industry

fats are used as fertilizers, which are necessary for life

plant life. In the human body, N and P play an important role - nitrogen

is part of amino acids, which are an integral part of proteins, phosphorus in

form Ca5 [(PO4 )3 OH] is part of bones. In the human body there are

lasts on average about 1.8 kg N.

Some characteristics of the atoms of the VA-group elements are given in

The most important characteristics of atoms of elements of the VA group

	Electrical
	negative
	ness (according to
atom, nm	Polling)
		increase in the number of electric
		throne layers;
		increase in atomic size;
		increase in atomic size;
		decrease in ion energy

		decrease in electronegative
		value;
For comparison, the electronegativity of H is 2.2; O – 3.44.

Nitrogen differs from other elements of the subgroup in its very small orbital

tal radius and high electronegativity, N – third in electrical

triple-negativity element, after F and O.

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Valence electrons N –2s2 2p3.

N 2s

Nitrogen, like other elements of the second period,

differs markedly from the elements of its subgroup:

the N atom has only 4 valence orbitals and in compounds can form

call only 4 covalent bonds;

due to its very small atomic radius, nitrogen forms very strong

a simple substance in any state of aggregation consists of individual

very strong N molecules

N and is highly inert;

in electronegativity N is second only to F and O;

nitrogen exhibits all possible oxidation states: -3, –2, -1, 0, +1, +2, +3, +4, +5.

The large number of oxidation states and variety of compounds makes

The chemistry of nitrogen is very complex. The complexity is also aggravated by the kinetic difficulties characteristic of many redox reactions.

differences due to very strong multiple bonds between atoms

N and N and O atoms. Therefore, electrode potentials are of little help in determining

division of OVR products.

The most stable compound N is a simple substance.

In aqueous solutions, especially acidic ones, the NH4 + ion is very stable.

Nitrogen is a component of air, from which N2 is obtained.

The main amount of N2 is used for the synthesis of ammonia, from which other nitrogen compounds are then obtained. Among nitrogen compounds, ammonia, nitric acid and their salts find the widest practical application..

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Annual world production of NH3 is ~ 97 million tons/year, nitrogen dioxide

slots – 27 million tons/year. The chemistry of these important N compounds will be discussed

Ren first of all, after discussing the properties of simple matter.

Simple substance

The N2 molecule is the strongest of all diatomic molecules of simple substances. Three common electron pairs in the N N molecule are located in bonded

calling orbitals, there are no electrons in antibonding orbitals - this is calling

leads to a very high chemical bond energy – 944 kJ/mol (for comparison

However, the binding energy in the O2 molecule is 495 kJ/mol). Strong bonding causes high inertness of molecular nitrogen. The name of this element is associated with the chemical inertness of nitrogen. In Greek, “nitrogen” means

says "lifeless".

Under normal conditions, N2 is a colorless, odorless and tasteless gas.

The boiling and melting points of N2 are close: –196О С, and –210О С.

Nitrogen is obtained by fractional distillation of air , – for this air

At low temperatures they liquefy and then begin to increase the temperature.

Of the air components, nitrogen has the lowest boiling point and

forms the lightest boiling fraction. In fractional distillation, one

temporarily receive oxygen and inert gases.

The main amount of N2 goes to the production of ammonia, in addition,

nitrogen is used to create an inert atmosphere, including during production

properties of some metals; liquid nitrogen is also used as a coolant

giving agent in the laboratory and in industry.

At room temperature, nitrogen reacts slowly only with Li to form

formation of Li3 N. When magnesium burns in air, together with MgO oxide it forms

Mg3 N2 is also present.

Nitrides. Binary compounds of nitrogen with elements less electrically

triple negative than N are called nitrides.

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Ionic nitrides contain N3– anion. Ionic nitrides form Li,

metals of II and IB groups; in aqueous solutions they undergo irreversible

mu hydrolysis.

Mg3 N2 + 6H2 O = 2NH3 + 3 Mg(OH)2

With p-block metals and some light nonmetals, nitrogen is

forms covalent nitrides, for example, AlN, BN.

Most d-metals form nonstoichiometric interstitial products with nitrogen at high temperatures, in which N atoms occupy empty

totes in crystal lattices of metals. Therefore, such nitrides in external

In appearance, they resemble metals in electrical and thermal conductivity, but differ

They are characterized by high chemical inertness, hardness and refractoriness.

For example, non-stoichiometric nitrides Ta and Ti melt at temperatures above 3200o C.

Nitrogen does not react directly with halogens, but interacts with oxygen only under extreme conditions(with electric

rank).

The most important from a practical standpoint is the reaction of nitrogen with H2, which produces ammonia.

N 2 + 3H 2  2NH 3 ; H0 = –92 kJ/mol.

The exothermic nature of this reaction indicates that the total bond strength in ammonia molecules is higher than in the original molecules. An increase in temperature in accordance with Le Chatelier's principle leads to a shift in equilibrium towards the endothermic reaction, i.e. in the direction of ammonia decomposition. However, under normal conditions the reaction is extremely slow.

but the activation energy required to weaken the strong bonds in nitrogen and hydrogen molecules is too high. Therefore, the process must be carried out at a temperature of about 5000 C. To shift the equilibrium at high temperatures to the right, the pressure is increased to 300 - 500 atm, while the equilibrium

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This shifts in the direction of the reaction that occurs with a decrease in the number of gas molecules, i.e. in the direction of ammonia formation. Increased speed is achieved through the use of catalysts. A fused catalyst based on

new Fe3 O4 with additions of Al2 O3 and SiO2 and a catalyst based on metal

Fe. The synthesis of ammonia from nitrogen and hydrogen is the most important reaction in the pro-

industrial chemistry of nitrogen.

Nitrogen compounds

Ammonia and ammonium salts

Nitrogen in ammonia and ammonium salts is in the minimum oxidation state (–3). The oxidation state (–3) is quite stable for nitrogen.

Ammonia under normal conditions is a colorless gas with the characteristic

strong pungent odor, familiar from the smell of “ammonia” (10% dis-

ammonia solution in water). This gas is lighter than air, so it can be collected in containers turned upside down. Ammonia easily liquefies. To do this, it is enough to cool it at normal pressure to –33.5o C. The same effect

This effect can be achieved at room temperature, but by increasing the pressure to

7 – 8 atm. At elevated pressure, liquid ammonia is stored in steel containers.

nah. As liquid ammonia evaporates, it causes cooling in the environment. This is the basis for its use in refrigeration technology. The easy liquefaction of ammonia is due to hydrogen bonds between its molecules. The strength of hydrogen bonds between ammonia molecules is due to the very high electronegativity of nitrogen.

Liquid ammonia is colorless and undergoes autoprotolysis:

2NH3  NH4 + + NH2 –

The constant of this equilibrium is 2. 10–23 (at –50o C). Liquid ammonia

is a good ionizing solvent . Ammonium salts and weak

acids, for example, H2 S, dissolved in liquid ammonia become strong

with acids.

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Ammonia is highly soluble in water. The high solubility of ammonia in water (up to 700 volumes of NH3 in one volume of water) also explains the formation

We eat hydrogen bonds, but with water molecules. Concentrated dis-

the solution contains 25 mass% ammonia and has a density of 0.91 g/cm3. The molar concentration of NH3 in concentrated aqueous solutions reaches ~13

The NH3 molecule has a pyramidal structure, which is explained by sp3 -

hybridization of nitrogen valence atomic orbitals. One of the vertices of the tetrahedron

ra is occupied by a lone pair of electrons. The N–H bond is quite strong,

the bond energy is 389 kJ/mol, the bond length is 0.1 nm, the angle between bonds

north –108.3o. Upon addition of the H+ cation due to the unshared electron-

pair N, a tetrahedral very stable ammonium ion is formed

NH4+.

The presence of a lone electron pair at N in the NH3 molecule causes

celebrates many properties characteristic of ammonia.

The NH3 molecule is a good electron pair donor (DEP),

those. Lewis base, and very good proton acceptor A(H+),

those. Bronsted basis:

NH3 + H+  NH4 + . NH3 accepts a proton, like OH– ions: OH– + H+  H2 O

The acceptor properties of NH3 are weaker than those of the OH– anion. The protolysis constant for NH3 is 1.8. 109, and for the OH– ion – 1014.

Reactions with acids are the most characteristic reactions for NH3.

The ability of ammonia to form donor-acceptor bonds

so great that it can tear off hydrogen ions from such a strong bond

unity like water.

NH3 + H–– OH  NH4 + ), and the amount of NH4 + and OH– products is small compared to the equilibrium ammonia concentration. Aqueous solutions of ammonia behave like weak bases. According to established tradition, ammonia is often designated

have the formula NH4 OH and are called ammonium hydroxide, but the molecules

There is no NH4 OH in solution. The alkaline reaction of an aqueous NH3 solution is often described

is not described as the above equilibrium, but as the dissociation of molecules

NH4OH:

NH4 OH NH4 + + OH–

The constant of this equilibrium is 1.8. 10–5. In one liter one-molar

of ammonia solution, the concentration of NH4 + and OH– ions is 3.9. 10–3

mol/l, pH = 11.6.

The equilibrium between ammonia and OH– can be strongly shifted to the right by cations of some metals, which form insoluble hydroxides with OH– ions.

FeCl3 + 3NH3 + 3H–OH  Fe(OH)3  + 3NH4 Cl.

Ammonia can be used to produce insoluble bases.

When acids act on aqueous solutions of ammonia, ammonium salts are formed.

NH3 + HCl = NH4Cl

Almost all ammonium salts are colorless and soluble in water.

The equilibrium NH3 + H+  NH4 + is strongly shifted to the right (K = 1.8.109),

this means that NH3 is a strong proton acceptor, and the NH4+ cation

is a weak H donor+ , i.e. Bronsted acid. When alkali is added to ammonium salts, ammonia is formed, which is easily determined by the

NH4 Cl + NaOH = NH3 + H2 O + NaCl.

This reaction is usually used to detect ammonium ions in solution.

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Similar reactions can be used for laboratory production

NH3.

Ammonium chloride (called ammonia) reacts with oxides on metal surfaces like an acid at high temperatures, exposing pure metal. This is also the basis for the use of solid salt NH4 Cl when soldering metals. The “acidic” H+ from the NH4+ ion is capable of oxidizing very reactive metals, such as Mg.

Mg + 2NH4 Cl = H2 + MgCl2 + 2NH3

A characteristic property of ammonium salts is their thermal instability

tenacity. When heated, they decompose quite easily. Products

positions are determined by the properties of the acid anion. If the anion exhibits oxidizing properties, then NH4 + is oxidized and the oxidizing anion is reduced.

NH4 NO2 = N2 + 2H2 O

NH4 NO3 = N2 O + 2H2 O or 2NH4 NO3 = N2 + O2 + 4H2 O

(NH4 )2 Cr2 O7 = N2 + Cr2 O3 + 4H2 O

Ammonia and acid (or its anhydride) are released from the salts of volatile acids.

read), and in the case of non-volatile acids (for example, H3 PO4) - only NH3. NH4 HCO3 = NH3 + H2 O + CO2

Ammonium bicarbonate NH4 HCO3 is used in baking

Industrially, the resulting gases give the dough the necessary porosity.

Ammonium salts are used in the production of explosives and in

as nitrogen fertilizers. Ammonal, used in blasting practice, is a mixture of NH4 NO3 salt (72%), Al powder (25%) and carbon

la (3%). This mixture explodes only after detonation.

The second type of reactions in which NH3 exhibits the properties of an electron donor

throne pair is formation of amine complexes. Ammonia as a ligand attaches to cations of many d-elements, forming chemical

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NITROGEN, N (nitrogenium) , chemical element (at. number 7) VA subgroup of the periodic table of elements. The Earth's atmosphere contains 78% (vol.) nitrogen. To show how large these reserves of nitrogen are, we note that in the atmosphere above each square kilometer of the earth's surface there is so much nitrogen that up to 50 million tons of sodium nitrate or 10 million tons of ammonia (a compound of nitrogen with hydrogen) can be obtained from it, and yet this constitutes a small fraction of the nitrogen contained in the earth's crust. The existence of free nitrogen indicates its inertness and the difficulty of interacting with other elements at ordinary temperatures. Fixed nitrogen is part of both organic and inorganic matter. Plant and animal life contain nitrogen bound to carbon and oxygen in proteins. In addition, nitrogen-containing inorganic compounds such as nitrates (NO 3), nitrites (NO2), cyanides (CN), nitrides (N3) and azides (N3). Historical information. The experiments of A. Lavoisier, devoted to the study of the role of the atmosphere in maintaining life and combustion processes, confirmed the existence of a relatively inert substance in the atmosphere. Without establishing the elemental nature of the gas remaining after combustion, Lavoisier called it azote, which means “lifeless” in ancient Greek. In 1772, D. Rutherford from Edinburgh established that this gas is an element and called it “harmful air.” The Latin name for nitrogen comes from the Greek words nitron and gen, which means "saltpeter-forming".Nitrogen fixation and the nitrogen cycle. The term "nitrogen fixation" refers to the process of fixing atmospheric nitrogen N 2 . In nature, this can happen in two ways: either leguminous plants, such as peas, clover and soybeans, accumulate nodules on their roots, in which nitrogen-fixing bacteria convert it into nitrates, or atmospheric nitrogen is oxidized by oxygen under lightning conditions. S. Arrhenius found that up to 400 million tons of nitrogen are fixed annually in this way. In the atmosphere, nitrogen oxides combine with rainwater to form nitric and nitrous acids. In addition, it has been established that with rain and snow, approx. 6700 g nitrogen; reaching the soil, they turn into nitrites and nitrates. Plants use nitrates to form plant proteins. Animals, feeding on these plants, assimilate the protein substances of the plants and convert them into animal proteins. After the death of animals and plants, they decompose and nitrogen compounds turn into ammonia. Ammonia is used in two ways: bacteria that do not form nitrates break it down to elements, releasing nitrogen and hydrogen, and other bacteria form nitrites from it, which are oxidized by other bacteria to nitrates. This is how the nitrogen cycle occurs in nature, or the nitrogen cycle.Structure of the nucleus and electron shells. There are two stable isotopes of nitrogen in nature: with a mass number of 14 (contains 7 protons and 7 neutrons) and withmass number 15 (contains 7 protons and 8 neutrons). Their ratio is 99.635:0.365, so the atomic mass of nitrogen is 14.008. Unstable isotopes of nitrogen 12 N, 13 N, 16 N, 17 N are obtained artificially.Schematic electronic structure of the nitrogen atom is: 1 s 2 2s 2 2p x 1 2p y 1 2p z 1 . Consequently, the outer (second) electron shell contains 5 electrons that can participate in the formation of chemical bonds; nitrogen orbitals can also accept electrons, i.e. it is possible to form compounds with oxidation states from ( II I) to (V), and they are known. See also ATOMIC STRUCTURE.Molecular nitrogen. From determinations of gas density it has been established that the nitrogen molecule is diatomic, i.e. the molecular formula of nitrogen is Nє N (or N 2 ). Two nitrogen atoms have three outer ones 2p- electrons of each atom form a triple bond:N:::N:, forming electron pairs. Measured interatomic distance N N equals 1.095 Å . As is the case with hydrogen (cm. HYDROGEN), there are nitrogen molecules with different nuclear spins - symmetric and antisymmetric. At ordinary temperatures, the ratio of symmetric and antisymmetric forms is 2:1. In the solid state, two modifications of nitrogen are known: a cubic and b hexagonal with transition temperature a® b 237.39° C. Modification b melts at 209.96° C and boils at 195.78° C at 1 atm ( cm. table 1). Dissociation energy of a mole (28.016 g or 6.023 Ch 10 23 molecules) molecular nitrogen into atoms ( N 2 2N) equal to approximately 225 kcal. Therefore, atomic nitrogen can be formed during a quiet electrical discharge and is chemically more active than molecular nitrogen.Receipt and application. The method of obtaining elemental nitrogen depends on the required purity. Nitrogen is obtained in huge quantities for the synthesis of ammonia, while small admixtures of noble gases are acceptable.Nitrogen from the atmosphere. Economically, the release of nitrogen from the atmosphere is due to the low cost method of liquefying purified air (water vapor, CO 2 , dust and other impurities are removed). Successive cycles of compression, cooling and expansion of such air lead to its liquefaction. Liquid air is subjected to fractional distillation with a slow rise in temperature. The noble gases are released first, then nitrogen, and liquid oxygen remains. Purification is achieved by repeated fractionation processes. This method produces many millions of tons of nitrogen annually, mainly for the synthesis of ammonia, which is the feedstock in the production technology of various nitrogen-containing compounds for industry and agriculture. In addition, a purified nitrogen atmosphere is often used when the presence of oxygen is unacceptable.Laboratory methods. Nitrogen can be obtained in small quantities in the laboratory in various ways by oxidizing ammonia or ammonium ion, for example:

The process of oxidation of ammonium ion with nitrite ion is very convenient:Other methods are also known: decomposition of azides when heated, decomposition of ammonia with copper(II) oxide, interaction of nitrites with sulfamic acid or urea:

The catalytic decomposition of ammonia at high temperatures can also produce nitrogen: Physical properties. Some physical properties of nitrogen are given in table. 1.

Table 1. SOME PHYSICAL PROPERTIES OF NITROGEN
Density, g/cm 3	0.808 (liquid)
Melting point, °C	–209,96
Boiling point, °C	–195,8
Critical temperature, °C	–147,1
Critical pressure, atm a	33,5
Critical density, g/cm 3 a	0,311
Specific heat capacity, J/(mol K)	14.56 (15° C)
Electronegativity according to Pauling	3
Covalent radius,	0,74
Crystal radius,	1.4 (M 3–)
Ionization potential, V b
first	14,54
second	29,60
A Temperature and pressure at which densitiesNitrogen liquid and gaseous states are the same. b The amount of energy required to remove the first outer electron and the next one, per 1 mole of atomic nitrogen.

Chemical properties. As already noted, the predominant property of nitrogen under normal conditions of temperature and pressure is its inertness, or low chemical activity. The electronic structure of nitrogen contains an electron pair of 2 s-level and three half filled 2 r-orbitals, so one nitrogen atom can bind no more than four other atoms, i.e. its coordination number is four. The small size of an atom also limits the number of atoms or groups of atoms that can be associated with it. Therefore, many compounds of other members of the VA subgroup either have no analogues among nitrogen compounds, or similar nitrogen compounds turn out to be unstable. So, PCl 5 stable compound, and 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 xN y . In these compounds, nitrogen is negatively charged. In table Table 2 shows the oxidation states and examples of corresponding compounds.

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

Nitrides. Compounds of nitrogen with more electropositive elements, metals and nonmetals nitrides are similar to carbides and hydrides. They can be divided depending on the nature of the MN bond into ionic, covalent and with an intermediate type of bond. As a rule, these are crystalline substances.Ionic nitrides. The bonding in these compounds involves the transfer of electrons from the metal to nitrogen to form the N ion 3 . Such nitrides include Li 3 N, Mg 3 N 2, Zn 3 N 2 and Cu 3 N 2 . Apart from lithium, other alkali metals do not form IA subgroups of nitrides. Ionic nitrides have high melting points and react with water to form NH 3 and metal hydroxides.Covalent nitrides. When nitrogen electrons participate in the formation of a bond together with the electrons of another element without transferring them from nitrogen to another atom, nitrides with a covalent bond are formed. Hydrogen nitrides (such as ammonia and hydrazine) are completely covalent, as are nitrogen halides (NF 3 and NCl 3 ). Covalent nitrides include, for example, Si 3 N 4 , P 3 N 5 and BN are highly stable white substances, and BN has two allotropic modifications: hexagonal and diamond-like. The latter is formed at high pressures and temperatures and has a hardness close to that of diamond.Nitrides with an intermediate type of bond. Transition elements in reaction with NH 3 at high temperatures form an unusual class of compounds in which nitrogen atoms are distributed between regularly spaced metal atoms. There is no clear electron displacement in these compounds. Examples of such nitrides Fe 4 N, W 2 N, Mo 2 N, Mn 3 N 2 . These compounds are usually completely inert and have good electrical conductivity.Hydrogen compounds of nitrogen. Nitrogen and hydrogen interact to form compounds vaguely reminiscent of hydrocarbons (see also ORGANIC CHEMISTRY). The stability of hydrogen nitrates decreases with increasing number of nitrogen atoms in the chain, in contrast to hydrocarbons, which are stable in long chains. The most important hydrogen nitrides are ammonia NH 3 and hydrazine N 2 H 4 . These also include hydronitric acid HNNN (HN 3). Ammonia NH3. Ammonia is one of the most important industrial products of the modern economy. At the end of the 20th century. The USA produced approx. 13 million tons of ammonia annually (in terms of anhydrous ammonia).Molecule structure. NH 3 molecule has an almost pyramidal structure. The bond angle HNH is 107° , which is close to the tetrahedral angle 109° . The lone electron pair is equivalent to the attached group, resulting in the coordination number of nitrogen being 4 and the nitrogen being located at the center of the tetrahedron.Properties of ammonia. Some physical properties of ammonia in comparison with water are given in table. 3.

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

The boiling and melting points of ammonia are much lower than those of water, despite the similarity of molecular weights and the similarity of molecular structure. This is explained by the relatively greater strength of intermolecular bonds in water than in ammonia (such intermolecular bonds are called hydrogen bonds).Ammonia as a solvent. The high dielectric constant and dipole moment of liquid ammonia make it possible to use it as a solvent for polar or ionic inorganic substances. Ammonia solvent occupies an intermediate position between water and organic solvents such as ethyl alcohol. Alkali and alkaline earth metals dissolve in ammonia, forming dark blue solutions. It can be assumed that solvation and ionization of valence electrons occurs in solution according to the schemeThe blue color is associated with solvation and the movement of electrons or the mobility of “holes” in a liquid. At a high concentration of sodium in liquid ammonia, the solution takes on a bronze color and is highly electrically conductive. Unbound alkali metal can be separated from such a solution by evaporation of ammonia or the addition of sodium chloride. Solutions of metals in ammonia are good reducing agents. Autoionization occurs in liquid ammonia

similar to the process occurring in water:

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

Liquid ammonia as a solvent has an advantage in some cases where it is not possible to carry out reactions in water due to the rapid interaction of components with water (for example, oxidation and reduction). For example, in liquid ammonia, calcium reacts with KCl to form CaCl 2 and K, since CaCl 2 is insoluble in liquid ammonia, and K is soluble, and the reaction proceeds completely. In water, such a reaction is impossible due to the rapid interaction of Ca with water.

Production of ammonia. Gaseous NH 3 is released from ammonium salts under the action of a strong base, for example, NaOH:The method is applicable in laboratory conditions. Small ammonia production is also based on the hydrolysis of nitrides, such as Mg 3 N 2 , water. Calcium cyanamide CaCN 2 When interacting with water, it also forms ammonia. The main industrial method for producing ammonia is its catalytic synthesis from atmospheric nitrogen and hydrogen at high temperature and pressure:

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
Aquatic environment	Ammonia environment
Neutralization
OH – + H 3 O + ® 2H 2 O	NH 2 – + NH 4 + ® 2NH 3
Hydrolysis (protolysis)


PCl 5 + 3H 2 O POCl 3 + 2H 3 O + + 2Cl –	PCl 5 + 4NH 3 PNCl 2 + 3NH 4 + + 3Cl –
Substitution
Zn + 2H 3 O + ® Zn 2+ + 2H 2 O + H 2	Zn + 2NH 4 + ® Zn 2+ + 2NH 3 + H 2
Solvation (complexation )
Al 2 Cl 6 + 12H 2 O 2 3+ + 6Cl –	Al 2 Cl 6 + 12NH 3 2 3+ + 6Cl –
Amphotericity
Zn 2+ + 2OH – Zn(OH) 2	Zn 2+ + 2NH 2 – Zn(NH 2) 2
Zn(OH) 2 + 2H 3 O + Zn 2+ + 4H 2 O	Zn(NH 2) 2 + 2NH 4 + Zn 2+ + 4NH 3
Zn(OH) 2 + 2OH – Zn(OH) 4 2–	Zn(NH 2) 2 + 2NH 2 – Zn(NH 2) 4 2–

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

Symbol M n+ represents a transition metal ion (B-subgroups of the periodic table, for example, Cu 2+, Mn 2+ andetc.). Any protic (i.e. H-containing) acid reacts with ammonia in aqueous solution to form ammonium salts, such as ammonium nitrate NH 4 NO 3 , ammonium chloride NH 4 Cl, ammonium sulfate (NH 4) 2 SO 4 , ammonium phosphate (NH 4) 3PO 4 . These salts are widely used in agriculture as fertilizers to introduce nitrogen into the soil. Ammonium nitrate is also used as an inexpensive explosive; it was first used with petroleum fuel (diesel oil). An aqueous solution of ammonia is used directly for introduction into the soil or with irrigation water. Urea NH 2 CONH 2 , obtained by synthesis from ammonia and carbon dioxide, is also a fertilizer. Ammonia gas reacts with metals such as Na and K to form amides:Ammonia also reacts with hydrides and nitrides to form amides:

Alkali metal amides (e.g. NaNH 2) react with N 2 O when heated, forming azides: Gaseous NH 3 reduces heavy metal oxides to metals at high temperatures, apparently due to hydrogen generated from the decomposition of ammonia into N 2 and H 2: Hydrogen atoms in the NH molecule 3 can be replaced by halogen. Iodine reacts with concentrated NH solution 3 , forming a mixture of substances containing N I 3 . This substance is very unstable and explodes at the slightest mechanical impact. When reacting NH 3 c Cl 2 chloramines NCl 3, NHCl 2 and NH 2 Cl are formed. When ammonia is exposed to sodium hypochlorite NaOCl (formed from NaOH and Cl2 ) the final product is hydrazine:

Hydrazine. The above reactions are a method for preparing hydrazine monohydrate of composition N 2 H 4 H H 2 O. Anhydrous hydrazine is formed by special distillation of the monohydrate with BaO or other water-removing substances. The properties of hydrazine are slightly similar to hydrogen peroxide H 2 O 2 . Pure Anhydrous Hydrazine colorless hygroscopic liquid, boiling at 113.5°C ; dissolves well in water, forming a weak base In an acidic environment (H+ ) hydrazine forms soluble hydrazonium salts of the + X type . The ease with which hydrazine and some of its derivatives (such as methylhydrazine) react with oxygen allows it to be used as a component of liquid rocket fuel. Hydrazine and all its derivatives are highly toxic.Nitrogen oxides. In compounds with oxygen, nitrogen exhibits all oxidation states, forming oxides: N2 O, NO, N 2 O 3, NO 2 (N 2 O 4), N 2 O 5. There is scant information on the formation of nitrogen peroxides (NO 3, NO 4). Nitric oxide (I) N 2 O (dianitrogen monoxide) is obtained from the thermal dissociation of ammonium nitrate:The molecule has a linear structureO is fairly inert at room temperature, but at high temperatures it can support the combustion of easily oxidized materials. N 2 O, known as laughing gas, is used for mild anesthesia in medicine.Nitric oxide(II) NO colorless gas, is one of the products of the catalytic thermal dissociation of ammonia in the presence of oxygen:

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

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

and also by reactions:

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

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

N 2 O 5 is a good oxidizing agent, it reacts easily, sometimes violently, with metals and organic compounds and, in a pure state, explodes when heated. Probable structure. When the solution is evaporated, a white explosive is formed with the expected structure HON=NOH. Nitrous acid HNO2 is not exists in its pure form, but aqueous solutions of its low concentration are formed by adding sulfuric acid to barium nitrite:Nitrous acid is also formed when an equimolar mixture of NO and NO is dissolved 2 (or N 2 O 3 ) in water. Nitrous acid is slightly stronger than acetic acid. The oxidation state of nitrogen in it is +3 (its structure HON=O), those. it can be both an oxidizing agent and a reducing agent. Under the influence of reducing agents, it is usually restored to NO , and when interacting with oxidizing agents, it is oxidized to nitric acid.

The rate of dissolution of some substances, such as metals or iodide ion, in nitric acid depends on the concentration of nitrous acid present as an impurity. Salts of nitrous acid nitrites are highly soluble in water, except for silver nitrite.

NaNO2 used in the production of dyes.Nitric acid HNO3 one of the most important inorganic products of the main chemical industry. It is used in the technologies of many other inorganic and organic substances, such as explosives, fertilizers, polymers and fibers, dyes, pharmaceuticals, etc. See also CHEMICAL ELEMENTS.LITERATURE Nitrogenist's Directory. M., 1969
Nekrasov B.V. Basics of general chemistry. M., 1973
Nitrogen fixation problems. Inorganic and physical chemistry. M., 1982

Compounds with oxidation state –3. Nitrogen compounds in the -3 oxidation state are represented by ammonia and metal nitrides.

Ammonia- NH 3 is a colorless gas with a characteristic pungent odor. The ammonia molecule has the geometry of a trigonal pyramid with a nitrogen atom at the apex. The atomic orbitals of nitrogen are in sp 3- hybrid state. Three orbitals are involved in the formation of nitrogen-hydrogen bonds, and the fourth orbital contains a lone electron pair, the molecule has a pyramidal shape. The repulsive action of the lone pair of electrons causes the bond angle to decrease from the expected 109.5° to 107.3°.

At a temperature of -33.4 °C, ammonia condenses, forming a liquid with a very high heat of evaporation, which makes it possible to use it as a refrigerant in industrial refrigeration units.

The presence of a lone electron pair on the nitrogen atom allows it to form another covalent bond through the donor-acceptor mechanism. Thus, in an acidic environment, the formation of the molecular ammonium cation - NH 4 + occurs. The formation of the fourth covalent bond leads to the alignment of bond angles (109.5°) due to the uniform repulsion of hydrogen atoms.

Liquid ammonia is a good self-ionizing solvent:

2NH 3 NH 4 + + NH 2 -

amide anion

Alkali and alkaline earth metals dissolve in it, forming colored conductive solutions. In the presence of a catalyst (FeCl 3), the dissolved metal reacts with ammonia to release hydrogen and form an amide, for example:

2Na + 2NH 3 = 2NaNH 2 + H 2

sodium amide

Ammonia is very soluble in water (at 20 °C, about 700 volumes of ammonia dissolve in one volume of water). In aqueous solutions it exhibits the properties of a weak base.

NH 3 + H 2 O ® NH 3 ×H 2 O NH 4 + + OH -

= 1.85·10 -5

In an oxygen atmosphere, ammonia burns to form nitrogen; on a platinum catalyst, ammonia is oxidized to nitrogen oxide (II):

4NH 3 + 3O 2 = 2N 2 + 6H 2 O; 4NH 3 + 5O 2 = 4NO + 6H 2 O

As a base, ammonia reacts with acids to form ammonium cation salts, for example:

NH 3 + HCl = NH 4 Cl

Ammonium salts are highly soluble in water and slightly hydrolyzed. In the crystalline state they are thermally unstable. The composition of thermolysis products depends on the properties of the acid forming the salt:

NH 4 Cl ® NH 3 + HCl; (NH 4) 2 SO 4 ® NH 3 + (NH 4)HSO 4

(NH 4) 2 Cr 2 O 7 ® N 2 + Cr 2 O 3 + 4H 2 O

When aqueous solutions of ammonium salts are exposed to alkalis when heated, ammonia is released, which allows this reaction to be used as a qualitative reaction for ammonium salts and as a laboratory method for producing ammonia.

NH 4 Cl + NaOH = NaCl + NH 3 + H 2 O

In industry, ammonia is produced by direct synthesis.

N 2 + 3H 2 2NH 3

Since the reaction is highly reversible, the synthesis is carried out at elevated pressure (up to 100 mPa). To speed up the process, it is carried out in the presence of a catalyst (sponge iron promoted by additives) and at a temperature of about 500 °C.

Nitrides are formed as a result of reactions of many metals and non-metals with nitrogen. The properties of nitrides naturally change over time. For example, for elements of the third period:

Nitrides of s-elements of groups I and II are crystalline salt-like substances that easily decompose with water to form ammonia.

Li 3 N + 3H 2 O = 3LiOH + NH 3

Of the halogen nitrides in the free state, only Cl 3 N is isolated; the acidic character manifests itself in the reaction with water:

Cl 3 N + 3H 2 O = 3HClO + NH 3

The interaction of nitrides of different natures leads to the formation of mixed nitrides:

Li 3 N + AlN = Li 3 AlN 2; 5Li 3 N + Ge 3 N 4 = 3Li 5 GeN 3

lithium nitridegermanate(IV) nitridealuminate

Nitrides BN, AlN, Si 3 N 4, Ge 3 N 4 are solid polymer substances with high melting points (2000-3000 ° C), they are semiconductors or dielectrics. D-metal nitrides are crystalline compounds of variable composition (bertolides), very hard, refractory and chemically stable, exhibit metallic properties: metallic luster, electrical conductivity.

Compounds with oxidation state –2. Hydrazine - N 2 H 4 - the most important inorganic nitrogen compound in the oxidation state -2.

Hydrazine is a colorless liquid with a boiling point of 113.5 °C, fuming in air. Hydrazine vapors are extremely toxic and form explosive mixtures with air. Hydrazine is obtained by oxidizing ammonia with sodium hypochlorite:

2N -3 H 3 + NaCl +1 O = N 2 -2 H 4 + NaCl -1 + H 2 O

Hydrazine mixes with water in any ratio and in solution behaves as a weak diacid base, forming two series of salts.

N 2 H 4 + H 2 O N 2 H 5 + + OH - , K b = 9.3 × 10 -7 ;

hydrozonium cation

N 2 H 5 + + H 2 O N 2 H 6 2+ + OH - , K b = 8.5 × 10 -15 ;

dihydrosonium cation

N 2 H 4 + HCl N 2 H 5 Cl; N 2 H 5 Cl + HCl N 2 H 6 Cl 2

hydrozonium chloride dihydrosonium dichloride

Hydrazine is the strongest reducing agent:

4KMn +7 O 4 + 5N 2 -2 H 4 + 6H 2 SO 4 = 5N 2 0 + 4Mn +2 SO 4 + 2K 2 SO 4 + 16H 2 O

Unsymmetrical dimethylhydrazine (heptyl) is widely used as rocket fuel.

Compounds with oxidation state –1. Hydroxylamine - NH 2 OH - is the main inorganic nitrogen compound in the oxidation state -1.

Hydroxylamine is obtained by reducing nitric acid with hydrogen at the time of release during electrolysis:

HNO 3 + 6H = NH 2 OH + 2H 2 O

This is a colorless crystalline substance (mp 33 °C), highly soluble in water, in which it exhibits the properties of a weak base. With acids it produces hydroxylammonium salts - stable, colorless substances soluble in water.

NH 2 OH + H 2 O + + OH - , K b = 2×10 -8

hydroxylammonium ion

The nitrogen atom in the NH 2 OH molecule exhibits an intermediate oxidation state (between -3 and +5), so hydroxylamine can act as both a reducing agent and an oxidizing agent:

2N -1 H 2 OH + I 2 + 2KOH = N 0 2 + 2KI + 4H 2 O;

reducing agent

2N -1 H 2 OH + 4FeSO 4 + 3H 2 SO 4 = 2Fe 2 (SO 4) 3 + (N -3 H 4) 2 SO 4 + 2H 2 O

oxidant

NH 2 OH easily decomposes when heated, undergoing disproportionation:

3N -1 H 2 OH = N 0 2 + N -3 H 3 + 3H 2 O;

Compounds with oxidation state +1. Nitric oxide (I) - N 2 O (nitrous oxide, laughing gas). The structure of its molecule can be conveyed by the resonance of two valence schemes, which show that this compound can only be considered formally as nitrogen(I) oxide; in reality it is nitrogen(V) oxonitride - ON +5 N -3.

N 2 O is a colorless gas with a faint pleasant odor. In small concentrations it causes bouts of unbridled joy, in large doses it has a general anesthetic effect. A mixture of nitrous oxide (80%) and oxygen (20%) was used in medicine for anesthesia.

In laboratory conditions, nitric oxide (I) can be obtained by the decomposition of ammonium nitrate. N 2 O obtained by this method contains impurities of higher nitrogen oxides, which are extremely toxic!

NH 4 NO 3 ¾® N 2 O + 2H 2 O

In terms of chemical properties, nitric oxide (I) is a typical non-salt-forming oxide; it does not react with water, acids and alkalis. When heated, it decomposes to form oxygen and nitrogen. For this reason, N 2 O can act as an oxidizing agent, for example:

N 2 O + H 2 = N 2 + H 2 O

Compounds with oxidation state +2. Nitrogen(II) oxide - NO - a colorless gas, extremely toxic. In air it is quickly oxidized by oxygen to form no less toxic nitrogen oxide (IV). In industry, NO is produced by the oxidation of ammonia on a platinum catalyst or by passing air through an electric arc (3000-4000 °C).

4NH 3 + 5O 2 = 4NO + 6H 2 O; N2 + O2 = 2NO

A laboratory method for producing nitric oxide (II) is the reaction of copper with dilute nitric acid.

3Cu + 8HNO 3 (diluted) = 3Cu(NO 3) 2 + 2NO + 4H 2 O

Nitrogen(II) oxide is a non-salt-forming oxide, a strong reducing agent, and easily reacts with oxygen and halogens.

2NO + O 2 = 2NO 2; 2NO + Cl 2 = 2NOCl

nitrosyl chloride

At the same time, when interacting with strong reducing agents, NO acts as an oxidizing agent:

2NO + 2H 2 = N 2 + 2H 2 O; 10NO + 4P = 5N 2 + 2P 2 O 5

Compounds with oxidation state +3. Nitrogen(III) oxide - N 2 O 3 - liquid of intense blue color (temperature -100 °C). Stable only in liquid and solid states at low temperatures. Apparently exists in two forms:

Nitrogen(III) oxide is obtained by joint condensation of NO and NO 2 vapors. Dissociates in liquids and vapors.

NO 2 + NO N 2 O 3

The properties are typical acid oxide. Reacts with water, forming nitrous acid, and with alkalis it forms salts - nitrites.

N 2 O 3 + H 2 O = 2HNO 2; N 2 O 3 + 2NaOH = 2NaNO 2 + H 2 O

Nitrous acid- medium strength acid (K a = 1×10 -4). It is not isolated in its pure form; in solutions it exists in two tautomeric forms (tautomers are isomers that are in dynamic equilibrium).

nitrite form nitro form

Nitrous acid salts are stable. The nitrite anion exhibits pronounced redox duality. Depending on the conditions, it can perform both the function of an oxidizing agent and the function of a reducing agent, for example:

2NaNO 2 + 2KI + 2H 2 SO 4 = I 2 + 2NO + K 2 SO 4 + Na 2 SO 4 + 2H 2 O

oxidant

KMnO 4 + 5NaNO 2 + 3H 2 SO 4 = 2MnSO 4 + 5NaNO 3 + K 2 SO 4 + 3H 2 O

reducing agent

Nitrous acid and nitrites tend to disproportionate:

3HN +3 O 2 = HN +5 O 3 + 2N +2 O + H 2 O

Compounds with oxidation state +4. Nitrogen oxide (IV) - NO 2 - brown gas, with a pungent unpleasant odor. Extremely toxic! In industry, NO 2 is produced by the oxidation of NO. A laboratory method for producing NO 2 is the interaction of copper with concentrated nitric acid, as well as the thermal decomposition of lead nitrate.

Cu + 4HNO 3 (conc.) = Cu(NO 3) 2 + 2NO 2 + 2H 2 O;

2Pb(NO 3) 2 = 2PbO + 4NO 2 + O 2

The NO 2 molecule has one unpaired electron and is a stable free radical, so nitric oxide dimerizes easily.

The dimerization process is reversible and very sensitive to temperature:

paramagnetic, diamagnetic,

brown colorless

Nitrogen dioxide is an acidic oxide that reacts with water, forming a mixture of nitric and nitrous acid (mixed anhydride).

2NO 2 + H 2 O = HNO 2 + HNO 3; 2NO 2 + 2NaOH = NaNO 3 + NaNO 2 + H 2 O

Compounds with oxidation state +5. Nitric oxide (V) - N 2 O 5 - a white crystalline substance. It is obtained by dehydration of nitric acid or oxidation of nitric oxide (IV) with ozone:

2HNO 3 + P 2 O 5 = N 2 O 5 + 2HPO 3; 2NO 2 + O 3 = N 2 O 5 + O 2

In the crystalline state, N 2 O 5 has a salt-like structure - + -, in vapors (sublime temperature 33 ° C) - molecular.

N 2 O 5 - acid oxide - nitric acid anhydride:

N2O5 + H2O = 2HNO3

Nitric acid- HNO 3 is a colorless liquid with a boiling point of 84.1 ° C, decomposes when heated and exposed to light.

4HNO 3 = 4NO 2 + O 2 + 2H 2 O

Impurities of nitrogen dioxide give concentrated nitric acid a yellow-brown color. Nitric acid mixes with water in any ratio and is one of the strongest mineral acids; it completely dissociates in solution.

The structure of the nitric acid molecule is described by the following structural formulas:

Difficulties in writing the structural formula of HNO 3 are caused by the fact that, exhibiting an oxidation state of +5 in this compound, nitrogen, as an element of the second period, can form only four covalent bonds.

Nitric acid is one of the strongest oxidizing agents. The depth of its recovery depends on many factors: concentration, temperature, reducing agent. Typically, oxidation with nitric acid produces a mixture of reduction products:

HN +5 O 3 ® N +4 O 2 ® N +2 O ® N +1 2 O ® N 0 2 ® +

The predominant product of oxidation of nonmetals and inactive metals with concentrated nitric acid is nitric oxide (IV):

I 2 + 10HNO 3 (conc) = 2HIO 3 + 10NO 2 + 4H 2 O;

Pb + 4HNO 3 (conc) = Pb(NO 3) 2 + 2NO 2 + 2H 2 O

Concentrated nitric acid passivates iron and aluminum. Aluminum is passivated even with dilute nitric acid. Nitric acid of any concentration has no effect on gold, platinum, tantalum, rhodium and iridium. Gold and platinum are dissolved in aqua regia - a mixture of concentrated nitric and hydrochloric acids in a ratio of 1: 3.

Au + HNO 3 + 4HCl = H + NO + 2H 2 O

The strong oxidizing effect of aqua regia is due to the formation of atomic chlorine during the decomposition of nitrosyl chloride, a product of the interaction of nitric acid with hydrogen chloride.

HNO 3 + 3HCl = Cl 2 + NOCl + 2H 2 O;

NOCl = NO + Cl×

An effective solvent for low-active metals is a mixture of concentrated nitric and hydrofluoric acids.

3Ta + 5HNO3 + 21HF = 3H2 + 5NO + 10H2O

Dilute nitric acid, when interacting with non-metals and low-active metals, is reduced predominantly to nitrogen oxide (II), for example:

3P + 5HNO 3 (dil) + 2H 2 O = 3H 3 PO 4 + 5NO;

3Pb + 8HNO 3 (dil) = 3Pb(NO 3) 2 + 2NO + 4H 2 O

Active metals reduce dilute nitric acid to N 2 O, N 2 or NH 4 NO 3, for example,

4Zn + 10HNO 3 (dil) = 4Zn(NO 3) 2 + NH 4 NO 3 + 3H 2 O

The bulk of nitric acid is used in the production of fertilizers and explosives.

Nitric acid is produced industrially by contact or arc methods, which differ in the first stage - the production of nitric oxide (II). The arc method is based on producing NO by passing air through an electric arc. In the contact method, NO is produced by the oxidation of ammonia with oxygen on a platinum catalyst. Next, nitrogen oxide(II) is oxidized to nitrogen oxide(IV) by atmospheric oxygen. By dissolving NO 2 in water in the presence of oxygen, nitric acid is obtained with a concentration of 60-65%.

4NO 2 + O 2 + 2H 2 O = 4HNO 3

If necessary, nitric acid is concentrated by distillation with concentrated sulfuric acid. In the laboratory, 100% nitric acid can be obtained by the action of concentrated sulfuric acid on crystalline sodium nitrate when heated.

NaNO 3 (cr) + H 2 SO 4 (conc) = HNO 3 + NaHSO 4

Nitric acid salts- nitrates - highly soluble in water, thermally unstable. The decomposition of nitrates of active metals (excluding lithium), located in the series of standard electrode potentials to the left of magnesium, leads to the formation of nitrites. For example:

2KNO 3 = 2KNO 2 + O 2

During the decomposition of lithium and magnesium nitrates, as well as metal nitrates located in the series of standard electrode potentials to the right of magnesium, up to copper, a mixture of nitrogen(IV) oxide and oxygen is released. For example:

2Cu(NO 3) 2 = 2CuO + 4NO 2 + O 2

Nitrates of metals located at the end of the activity series decompose to free metal:

2AgNO3 = 2Ag + 2NO2 + O2

Sodium, potassium and ammonium nitrates are widely used for the production of gunpowder and explosives, and also as nitrogen fertilizers (saltpeter). Ammonium sulfate, ammonia water and carbamide (urea) - a complete carbonic acid amide are also used as fertilizers:

Hydrogen azide(dinitridonitrate) - HN 3 (HNN 2) – colorless volatile liquid (melting point –80 °C, boiling point 37 °C) with a pungent odor. The central nitrogen atom is in sp-hybridization, the oxidation state is +5, the atoms adjacent to it have an oxidation state of –3. Molecule structure:

An aqueous solution of HN 3 - hydronitric acid is close in strength to acetic acid, K a = 2.6 × 10 -5. Stable in dilute solutions. It is obtained by reacting hydrazine and nitrous acid:

N 2 H 4 + HNO 2 = HN 3 + 2 H 2 O

The oxidative properties of HN 3 (HN +5 N 2) resemble nitric acid. So, if the interaction of a metal with nitric acid produces nitrogen oxide (II) and water, then with hydronitrous acid nitrogen and ammonia are formed. For example,

Cu + 3HN +5 N 2 = Cu(N 3) 2 + N 2 0 + NH 3

A mixture of HN 3 and HCl behaves like aqua regia. Salts of hydronitric acid - azides. Only alkali metal azides are relatively stable; at temperatures > 300 °C they destroy without explosion. The rest disintegrate explosively when struck or heated. Lead azide is used in the production of detonators:

Pb(N 3) 2 = Pb + 3N 2 0

The starting product for the preparation of azides is NaN 3, which is formed as a result of the reaction of sodium amide and nitric oxide (I):

NaNH 2 + N 2 O = NaN 3 + H 2 O

4.2.Phosphorus

Phosphorus is represented in nature by one isotope - 31 P, the clarke of phosphorus is 0.05 mol.%. It is found in the form of phosphate minerals: Ca 3 (PO 4) 2 - phosphorite, Ca 5 (PO 4) 3 X (X = F,Cl,OH) - apatites. It is part of the bones and teeth of animals and humans, as well as the composition of nucleic acids (DNA and RNA) and adenosine phosphoric acids (ATP, ADP and AMP).

Phosphorus is obtained by reducing phosphorite with coke in the presence of silicon dioxide.

Ca 3 (PO 4) 2 + 3SiO 2 + 5C = 3CaSiO 3 + 2P + 5CO

A simple substance - phosphorus - forms several allotropic modifications, of which the main ones are white, red and black phosphorus. White phosphorus is formed by condensation of phosphorus vapor and is a white waxy substance (mp 44 °C), insoluble in water, soluble in some organic solvents. White phosphorus has a molecular structure and consists of tetrahedral P4 molecules.

The bond tension (P-P-P bond angle is only 60°) causes the high reactivity and toxicity of white phosphorus (lethal dose of about 0.1 g). Since white phosphorus is highly soluble in fats, milk cannot be used as an antidote for poisoning. In air, white phosphorus spontaneously ignites, so it is stored in hermetically sealed chemical containers under a layer of water.

Red phosphorus has a polymer structure. It is obtained by heating white phosphorus or irradiating it with light. Unlike white phosphorus, it is slightly reactive and non-toxic. However, residual amounts of white phosphorus can make red phosphorus toxic!

Black phosphorus is obtained by heating white phosphorus under a pressure of 120 thousand atm. It has a polymer structure, has semiconductor properties, is chemically stable and non-toxic.

Chemical properties. White phosphorus is spontaneously oxidized by atmospheric oxygen at room temperature (oxidation of red and black phosphorus occurs when heated). The reaction occurs in two stages and is accompanied by luminescence (chemiluminescence).

2P + 3O 2 = 2P 2 O 3; P 2 O 3 + O 2 = P 2 O 5

Phosphorus also interacts stepwise with sulfur and halogens.

2P + 3Cl 2 = 2PCl 3 ; PCl 3 + Cl 2 = PCl 5

When interacting with active metals, phosphorus acts as an oxidizing agent, forming phosphides - phosphorus compounds in the -3 oxidation state.

3Ca + 2P = Ca 3 P 2

Oxidizing acids (nitric and concentrated sulfuric acids) oxidize phosphorus to phosphoric acid.

P + 5HNO 3 (conc) = H 3 PO 4 + 5NO 2 + H 2 O

When boiled with alkali solutions, white phosphorus disproportionates:

4P 0 + 3KOH + 3H 2 O = P -3 H 3 + 3KH 2 P +1 O 2

phosphine potassium hypophosphite

There are chemical elements that exhibit different oxidation states, which allows the formation of a large number of compounds with certain properties during chemical reactions. Knowing the electronic structure of an atom, we can guess what substances will be formed.

The oxidation state of nitrogen can vary from -3 to +5, which indicates the variety of compounds based on it.

Element characteristics

Nitrogen belongs to the chemical elements located in group 15, in the second period in D.I. Mendeleev’s periodic system. It is assigned the serial number 7 and the abbreviated letter designation N. Under normal conditions, a relatively inert element; special conditions are required for reactions to occur.

It occurs in nature as a diatomic colorless gas of atmospheric air with a volume fraction of more than 75%. Contained in protein molecules, nucleic acids and nitrogen-containing substances of inorganic origin.

Atomic structure

To determine the oxidation state of nitrogen in compounds, it is necessary to know its nuclear structure and study the electron shells.

The natural element is represented by two stable isotopes, with their mass number 14 or 15. The first nucleus contains 7 neutron and 7 proton particles, and the second contains 1 more neutron particle.

There are artificial varieties of its atom with a mass of 12-13 and 16-17, which have unstable nuclei.

When studying the electronic structure of atomic nitrogen, it is clear that there are two electron shells (inner and outer). The 1s orbital contains one pair of electrons.

On the second outer shell there are only five negatively charged particles: two in the 2s-sub-level and three in the 2p-orbital. The valence energy level has no free cells, which indicates the impossibility of separating its electron pair. The 2p orbital is considered to be only half filled with electrons, which allows the addition of 3 negatively charged particles. In this case, the oxidation state of nitrogen is -3.

Taking into account the structure of the orbitals, we can conclude that this element with a coordination number of 4 is maximally bonded with only four other atoms. To form three bonds, exchange mechanism is used, another one is formed in a pre-nor-no-accept-tor way.

Oxidation states of nitrogen in different compounds

The maximum number of negative particles that its atom can attach is 3. In this case, its oxidation state appears equal to -3, inherent in compounds such as NH 3 or ammonia, NH 4 + or ammonium and Me 3 N 2 nitrides. The latter substances are formed with increasing temperature through the interaction of nitrogen with metal atoms.

The largest number of negatively charged particles that an element can give off is equal to 5.

Two nitrogen atoms are capable of combining with each other to form stable compounds with an oxidation state of -2. Such a bond is observed in N 2 H 4 or hydrazines, in azides of various metals or MeN 3. The nitrogen atom adds 2 electrons to the vacant orbitals.

There is an oxidation state of -1 when a given element receives only 1 negative particle. For example, in NH 2 OH or hydroxylamine it is negatively charged.

There are positive signs of the oxidation state of nitrogen, when electron particles are taken from the outer energy layer. They vary from +1 to +5.

The charge 1+ is present on nitrogen in N 2 O (monovalent oxide) and in sodium hyponitrite with the formula Na 2 N 2 O 2.

In NO (divalent oxide), the element gives up two electrons and becomes positively charged (+2).

There is an oxidation state of nitrogen 3 (in the compound NaNO 2 or nitride and also in trivalent oxide). In this case, 3 electrons are split off.

Charge +4 occurs in an oxide with valence IV or its dimer (N 2 O 4).

The positive sign of the oxidation state (+5) appears in N 2 O 5 or in pentavalent oxide, in nitric acid and its derivative salts.

Compounds of nitrogen and hydrogen

Natural substances based on the above two elements resemble organic hydrocarbons. Only hydrogen nitrates lose their stability as the amount of atomic nitrogen increases.

The most significant hydrogen compounds include molecules of ammonia, hydrazine and hydronitric acid. They are obtained by reacting hydrogen with nitrogen, and the latter substance also contains oxygen.

What is ammonia

It is also called hydrogen nitride, and its chemical formula is NH 3 with a mass of 17. Under conditions of normal temperature and pressure, ammonia has the form of a colorless gas with a pungent ammonia odor. It is 2 times less dense than air and easily dissolves in an aqueous environment due to the polar structure of its molecule. Refers to low-hazard substances.

In industrial quantities, ammonia is produced using catalytic synthesis from hydrogen and nitrogen molecules. There are laboratory methods for producing ammonium salts and sodium nitrite.

The structure of ammonia

The pyramidal molecule contains one nitrogen and 3 hydrogen atoms. They are located in relation to each other at an angle of 107 degrees. In a tetrahedron-shaped molecule, nitrogen is located in the center. Due to three unpaired p-electrons, it is connected by polar bonds of a covalent nature with 3 atomic hydrogens, which each have 1 s-electron. This is how an ammonia molecule is formed. In this case, nitrogen exhibits an oxidation state of -3.

This element still has an unshared pair of electrons at the outer level, which creates a covalent bond with a hydrogen ion that has a positive charge. One element is a donor of negatively charged particles, and the other is an acceptor. This is how the ammonium ion NH 4 + is formed.

What is ammonium

It is classified as a positively charged polyatomic ion or cation. Ammonium is also classified as a chemical substance that cannot exist in the form of a molecule. It consists of ammonia and hydrogen.

Ammonium with a positive charge in the presence of various anions with a negative sign is capable of forming ammonium salts, in which they behave like metals with valency I. Ammonium compounds are also synthesized with its participation.

Many ammonium salts exist in the form of crystalline, colorless substances that are readily soluble in water. If the compounds of the NH 4 + ion are formed by volatile acids, then under heating conditions they decompose with the release of gaseous substances. Their subsequent cooling leads to a reversible process.

The stability of such salts depends on the strength of the acids from which they are formed. Stable ammonium compounds correspond to a strong acidic residue. For example, stable ammonium chloride is produced from hydrochloric acid. At temperatures up to 25 degrees, such salt does not decompose, which cannot be said about ammonium carbonate. The latter compound is often used in cooking to rise dough, replacing baking soda.

Confectioners simply call ammonium carbonate ammonium. This salt is used by brewers to improve the fermentation of brewer's yeast.

A qualitative reaction for the detection of ammonium ions is the action of alkali metal hydroxides on its compounds. In the presence of NH 4 +, ammonia is released.

Chemical structure of ammonium

The configuration of its ion resembles a regular tetrahedron with nitrogen at the center. Hydrogen atoms are located at the vertices of the figure. To calculate the oxidation state of nitrogen in ammonium, you need to remember that the total charge of the cation is +1, and each hydrogen ion is missing one electron, and there are only 4 of them. The total hydrogen potential is +4. If we subtract the charge of all hydrogen ions from the charge of the cation, we get: +1 - (+4) = -3. This means that nitrogen has an oxidation state of -3. In this case, it adds three electrons.

What are nitrides

Nitrogen is able to combine with more electropositive atoms of metallic and non-metallic nature. As a result, compounds similar to hydrides and carbides are formed. Such nitrogen-containing substances are called nitrides. Between the metal and the nitrogen atom in compounds there are covalent, ionic and intermediate bonds. It is this characteristic that underlies their classification.

Covalent nitrides include compounds in which chemical bonds do not transfer electrons from atomic nitrogen, but form a common electron cloud together with negatively charged particles of other atoms.

Examples of such substances are hydrogen nitrides, such as ammonia and hydrazine molecules, as well as nitrogen halides, which include trichlorides, tribromides and trifluorides. Their common electron pair belongs equally to the two atoms.

Ionic nitrides include compounds with a chemical bond formed by the transition of electrons from the metal element to free levels of nitrogen. The molecules of such substances exhibit polarity. Nitrides have a nitrogen oxidation state of 3-. Accordingly, the total charge of the metal will be 3+.

Such compounds include nitrides of magnesium, lithium, zinc or copper, with the exception of alkali metals. They have a high melting point.

Nitrides with an intermediate bond include substances in which metal and nitrogen atoms are evenly distributed and there is no clear displacement of the electron cloud. Such inert compounds include nitrides of iron, molybdenum, manganese and tungsten.

Description of trivalent nitrogen oxide

It is also called an anhydride obtained from nitrous acid with the formula HNO 2. Taking into account the oxidation states of nitrogen (3+) and oxygen (2-) in the trioxide, the ratio of element atoms is 2 to 3 or N 2 O 3.

The liquid and gaseous forms of anhydride are very unstable compounds; they easily decompose into 2 different oxides with valence IV and II.