No acidic or basic oxide. Nitrogen oxides and their properties. Chemical properties of ammonia

Nitric oxide (II) (mon (o) nitric oxide, nitric oxide, nitrosyl radical) NO - non-salt-forming nitrogen oxide. It is a colorless gas that is poorly soluble in water. Liquefies with difficulty; in liquid and solid form it has a blue color.

The presence of an unpaired electron determines the tendency of NO to form weakly bound N 2 O 2 dimers. These are fragile compounds with ΔH ° dimerization = 17 kJ. Liquid nitric oxide (II) is 25% composed of N 2 O 2 molecules, and solid oxide entirely consists of them.

	30.0061 g / mol
Physical properties
Condition (st. Conv.)	colorless gas
Density	0.00134 (gas) g / cm³
Thermal properties
Melting temperature	−163.6 ° C
Boiling temperature	−151.7 ° C
Enthalpy of formation (st. Conv.)	81 kJ / mol
Chemical properties
Water solubility	0.01 g / 100 ml
Classification
Reg. CAS number

Receiving

Nitric oxide (II) is the only nitrogen oxide that can be obtained directly from free elements by combining nitrogen with oxygen at high temperatures (1200-1300 ° C) or in an electric discharge. In nature, it is formed in the atmosphere during lightning discharges:

N 2 + O 2 → 2NO - 180.9 kJ 2NO + O 2 → 2NO 2.

With a decrease in temperature, nitrogen oxide (II) decomposes into nitrogen and oxygen, but if the temperature drops sharply, then the oxide that did not have time to decompose exists for a long time: at low temperatures, the decomposition rate is low. This quenching is called "quenching" and is used in one of the methods for producing nitric acid.

In the laboratory, it is usually obtained by the interaction of 30% HNO 3 with certain metals, for example, with copper:

3Cu + 8HNO 3 (30%) → 3Cu (NO 3) 2 + 2NO + 4H 2 O.

Cleaner NO contaminated with impurities can be obtained by the reactions:

FeCl 2 + NaNO 2 + 2HCl → FeCl 3 + NaCl + NO + H 2 O; 2HNO 2 + 2HI → 2NO + I 2 ↓ + 2H 2 O.

The industrial method is based on the oxidation of ammonia at high temperature and pressure with the participation of Cr 2 O 3 (as catalysts):

4NH 3 + 5O 2 → 4NO + 6H 2 O.

Chemical properties

At room temperature and atmospheric pressure, NO oxidation by atmospheric oxygen occurs instantly:

2NO + O 2 → 2NO 2

For NO, reactions of addition of halogens with the formation of nitrosyl halides are also characteristic; in this reaction, NO exhibits the properties of a reducing agent:

2NO + Cl 2 → 2NOCl (nitrosyl chloride).

In the presence of stronger reducing agents, NO exhibits oxidizing properties:

2SO 2 + 2NO → 2SO 3 + N 2.

In water, NO is slightly soluble and does not react with it, being a non-salt-forming oxide.

Physiological action

Like all nitrogen oxides (except for N 2 O), NO is toxic and, if inhaled, affects the respiratory tract.

Over the past two decades, it has been established that this NO molecule has a wide spectrum of biological action, which can be conditionally divided into regulatory, protective and harmful. NO, being one of the messengers, is involved in the regulation of intra- and intercellular signaling systems. Nitric oxide produced by vascular endothelial cells is responsible for relaxation of vascular smooth muscles and their expansion (vasodilation), prevents platelet aggregation and adhesion of neurophils to the endothelium, and participates in various processes in the nervous, reproductive and immune systems. NO also has cytotoxic and cytostatic properties. The killer cells of the immune system use nitric oxide to kill bacteria and cancer cells. Diseases such as essential arterial hypertension, ischemic heart disease, myocardial infarction, primary pulmonary hypertension, bronchial asthma, neurotic depression, epilepsy, neurodegenerative diseases (Alzheimer's disease, Parkinson's disease), diabetes mellitus and others, impotence are associated with impaired NO biosynthesis and metabolism. ...

Nitric oxide can be synthesized in several ways. Plants use a non-enzymatic photochemical reaction between NO 2 and carotenoids. In animals, synthesis is carried out by the family of NO synthases (NOS). NOS enzymes are members of a heme-containing superfamily of enzymes called monooxygenases. Depending on structure and function, NOS can be divided into three groups: endothelial (eNOS), neuronal (nNOS), and inducible (iNOS). The active site of any of the NO synthases includes an iron porphyrin complex containing an axially coordinated cysteine ​​or methionine. Although all NOS isoforms catalyze the formation of NO, they are all products of different genes, each of which has its own characteristics, both in the mechanisms of action and localization, and in biological significance for the organism. Therefore, these isoforms are also commonly subdivided into constitutive (cNOS) and inducible (iNOS) nitric oxide synthase.

cNOS is constantly in the cytoplasm, depends on the concentration of calcium ions and calmodulin (a protein that is an intracellular mediator of calcium ion transport) and promotes the release of a small amount of NO for a short period in response to receptor stimulation. Inducible NOS appears in cells only after their induction by bacterial endotoxins and some mediators of inflammation, such as gamma-interferon, tumor necrosis factor, etc. The amount of NO formed under the influence of iNOS can vary and reach large amounts (nanomoles). At the same time, NO production persists for a longer time.

A characteristic feature of NO is the ability to quickly (in less than 5 seconds) diffuse through the membrane of the cell that synthesized it into the intercellular space and easily (without the participation of receptors) penetrate into target cells. Inside the cell, it activates some enzymes and inhibits others, thus participating in the regulation of cellular functions. Essentially, nitrogen monoxide is a local tissue hormone. NO plays a key role in suppressing the activity of bacterial and tumor cells by either blocking some of their iron-containing enzymes, or by damaging their cellular structures with nitric oxide or free radicals generated from nitric oxide. At the same time, superoxide accumulates in the focus of inflammation, which causes damage to proteins and lipids of cell membranes, which explains its cytotoxic effect on the target cell. Consequently, NO, excessively accumulating in the cell, can act in two ways: on the one hand, it can cause DNA damage and, on the other, it can have a pro-inflammatory effect.

Nitric oxide is able to initiate the formation of blood vessels. In the case of myocardial infarction, nitric oxide plays a positive role, as it induces new vascular growth, but in cancer, the same process causes the development of tumors, contributing to the nutrition and growth of cancer cells. On the other hand, this improves the delivery of nitric oxide to tumor cells. DNA damage caused by NO is one of the reasons for the development of

Oxides are called binary compounds of chemical elements with an oxygen atom, in which the oxidation state is 2-. Nitrogen, which has a lower electronegative value, forms various combinations with oxygen. These compounds belong to different classes of substances. Nitric oxide contains oxygen in an amount that sets the valence of the element N. It ranges from 1 to 5.

What are the oxides

There are about a dozen nitrogenous compounds containing the O-element. Of these, five are the most common: monovalent oxide, bivalent oxide, trivalent oxide, tetravalent oxide and pentavalent nitrogen oxide.

The rest of the connections are considered less common. These include tetravalent nitric oxide in the form of a dimer, unstable molecules of nitrilazide, nitrosylazide, trinitramide, and a nitrate radical.

Nitric oxide formulas

Below are the designations of the most significant compounds of element N.

This is primarily nitric oxide, the formula of which consists of two chemical signs - N and O. They are followed by indices, depending on the degree of oxidation of atoms.

• Monovalent nitrogen oxide has the formula N 2 O. In it, the N atom is charged +1.
• Nitrogen divalent oxide has the formula NO. In it, the N atom is charged +2.
• Nitrogen trivalent oxide has the formula N 2 O 3. In it, the N atom is charged +3.
• Tetravalent nitric oxide, whose formula is NO 2, has an atom charge of N +4.
• The pentavalent oxygen compound is designated N 2 O 5. In it, the N atom is charged +5.

Description of monovalent nitric oxide

It is also called dinitrogen, nitrous oxide and laughing gas. The latter name comes from an action associated with intoxication.

Nitric oxide with valence I under normal temperature conditions exists in the form of a non-combustible gas, colorless, which exhibits a pleasant sweetish taste and smell. Air is lighter than this compound. The oxide dissolves in an aqueous medium, ethanol, ethers and sulfuric acid.

Water, alkaline and acidic solutions are not capable of reacting with it, it does not form salts. It does not ignite, but it is able to support the combustion process.

Ammonia converts nitric oxide into azide (N3NH4).

When combined with molecules of ethers, chloroethane and cyclopropane, an explosive mixture is formed.

The usual conditions contribute to its inertia. The substance is reduced by heating.

Description of Divalent Nitric Oxide

It is also called monoxide, oxide or nitrosyl radical. Under normal temperature conditions, it is a colorless non-flammable gas, slightly soluble in an aqueous medium. It is oxidized by air, resulting in NO 2. Its liquid and solid form turns blue.

Nitric oxide can be a reducing agent in interactions with halogens. Their addition product is nitrosyl halide, which has the formula NOBr.

Sulfur dioxide and other strong reducing agents oxidize NO to form N 2 molecules.

Description of Trivalent Nitric Oxide

They are called nitrogen anhydride. In the normal state, it can be a liquid, with a blue color, and the standard parameters of the environment convert the oxide into the form of a colorless gas. It is stable only at low temperatures.

Molecules N 2 O 3 dissociate during heating with the release of mono - and divalent oxide.

It adds water as an anhydride to obtain nitrous acid, and forms salts in the form of nitrites with alkalis.

Description of tetravalent nitric oxide

In another way, it is called dioxide. It exists in the form of a brown-red gas that has a pungent odor and may also be a yellowish liquid.

Refers to acidic oxides, which have a well-developed chemical activity.

Its molecules oxidize non-metals to form oxygen-containing compounds and free nitrogen.

The dioxide interacts with tetravalent and hexavalent sulfur oxide. It turns out sulfuric acid. The method of its synthesis is called nitrous.

Nitric oxide can be dissolved in an aqueous medium. Nitric acid is the result of this reaction. This process is called disproportionation. The intermediate component is nitrous acid, which decomposes quickly.

If you dissolve tetravalent nitrogen oxide in alkali, then solutions of nitrates and nitrites are formed. You can use its liquid form to interact with a metal, then you get an anhydrous salt.

Description of pentavalent nitrogen oxide

It is also called dinitrogen pentoxide, nitronium nitrate, nitrile nitrate, or nitric anhydride.

It exists in the form of colorless crystals that are volatile and unstable. Their stability is observed at low temperatures. This structure is formed by nitrate and nitrite ions.

In gaseous form, the substance is in the form of NO 2 –O – NO 2 anhydride.

Pentavalent nitric oxide has acidic properties. It decomposes easily with the release of oxygen.

The substance reacts with water, resulting in nitric acid.

Alkalis dissolve the anhydride with the release of nitrate salts.

How are nitrogen oxides obtained

Nitrous oxide N 2 O is formed by careful heating of ammonium nitrate in dry form, but this method can be accompanied by an explosion.

The preferred method for producing monovalent oxide is the action of concentrated nitric acid on sulfamic acid. Heating is considered the main condition.

Nitrosyl, or NO, is a special nitrogen oxide produced by the interaction of N 2 and O 2 molecules. An important condition for this process is strong heating above 1000 ° C.

The natural production method is associated with lightning discharges in the atmospheric air. This oxide quickly combines with oxygen molecules to form dioxide.

The laboratory method for the synthesis of NO is associated with the reaction of metals and non-concentrated nitric acid. An example of such a reaction would be the interaction of copper with HNO 3.

Another way of forming nitrogen monoxide is the reaction of ferrous chloride with sodium nitrite and hydrochloric acid. The result of the process is ferric iron and sodium chloride, water and the oxide itself.

On an industrial scale, it is mined by oxidizing ammonia molecules during heating and under high pressure. The accelerator of the process is platinum or chromium trivalent oxide.

Dioxide, or NO 2, is obtained by the interaction of trivalent oxide arsenic with 50% nitric acid, which is applied dropwise to the surface of a solid reagent. A mixture of bivalent and tetravalent nitrogen oxides is formed.

If it is cooled to a temperature of -30 ° C, then nitrous anhydride, or N 2 O 3, is synthesized.

In powder form, it is obtained in the case of passing an electric current through its gaseous form.

If the starch powder is acted upon with nitric acid with a concentration of 50%, then bivalent and tetravalent nitrogen oxide, carbon dioxide gas and water are released. Subsequently, from the obtained first two compounds, the N 2 O 3 molecule is formed.

As a result of thermal decomposition of the lead nitro compound, free oxygen and lead oxide are released.

Anhydride, or N 2 O 5, is formed due to the elimination of a water molecule from an acid by nitric action of phosphorus pentavalent oxide.

Another way to synthesize it is by passing dry chlorine through anhydrous silver nitrate.

If nitrogen dioxide is acted upon by ozone molecules, then N 2 O 5 is formed.

The most important nitrogen oxides are presented in Table 1.

Nitric oxide (V) is a solid, other oxides are gaseous under normal conditions. Of these, nitric oxide (II) and nitrogen oxide (IV) are of the greatest practical importance. All nitrogen oxides are poisonous, with the exception of nitric oxide (I).

Nitric oxide (I) N 2 O. At room temperature, N 2 0 is a colorless gas (t pl = _ 91 ° C, t bale = -89 ° C), odorless, sweetish in taste, slightly soluble in water. When inhaled in small quantities, N 2 0 causes convulsive laughter, which is why it is called "laughing gas". Molecule N 2 0 is linear, low-polarity. By the method of valence bonds, its structure is described using two resonance structures:

The bond between nitrogen atoms (0.113 nm) is only slightly longer than the triple bond in the N 2 molecule (0.110 nm).

Nitric oxide (1) is obtained by thermal decomposition of ammonium nitrate at a temperature slightly above its melting point (170 ° C):

NH 4 NO 3 → N 2 0 + 2H 2 0

More pure N 2 0 is formed by coproportionation of nitrite and a salt of hydrazine or hydroxylamine:

NH 3 OHCI + NaN0 2 = N 2 O + 2H 2 0 + NaCl

Nitric oxide (II) NO- a colorless gas, slightly soluble in water and does not interact chemically with it. It easily combines with oxygen to form nitric oxide (IV):

2NO + O 2 → 2NO 2 + 113 kJ

Nitric oxide (II) is obtained in the laboratory by the action of dilute nitric acid (ρ = 1.2 g / cm 3, ω = 33%) on copper. The reaction equation is:

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

The NO molecule has an odd number of external electrons, therefore, it has one excess electron. The unsaturated nature of the NO molecule is manifested in its ability to complex with certain metal ions. Thus, when NO is passed into a FeSO 4 solution, the latter turns brown due to the formation of a compound of the composition SO 4. When heated, this fragile compound decomposes.

Nitric oxide (II) is a typical reducing agent. It discolors acidified potassium permanganate solution:

5NO + 3KMn0 4 + 2H 2 S0 4 = 2MnS0 4 + 3KN0 3 + Mn (N0 3) 2 + 2H 2 0

easily oxidized by oxygen:

2NO + 0 2 = 2N0 2

The process proceeds at a very high speed, since both reacting particles are radicals.

Due to the presence of one unpaired electron in the antibonding 2π-orbital, nitric oxide (II) is characterized by the processes of one-electron oxidation with the formation of a cation nitrosyl (nitrosonium) N0 + : N0 - e - = N0 +. In this case, the multiplicity of the N-O bond increases to three, and its energy increases from 627 (NO) to 1046 (NO +) kJ / mol. Nitrosyl derivatives are covalent nitrogen oxyhalides NOX (X is halogen), as well as ionic salts, for example, nitrosonium perchlorate, nitrosonium selenate (N0) 2 Se0 4. Nitrosonium hydrogen sulfate is obtained by passing sulfur dioxide through fuming nitric acid:

HN0 3 + S0 2 =

Other nitrosonium salts can also be obtained by reacting nitrogen oxides with concentrated acids, for example:

N 2 0 3 + H 2 Se0 4 = (N0) 2 Se0 4 + H 2 0

Nitrosonium salts are thermally unstable, and in the presence of water they irreversibly hydrolyze:

2 + H 2 0 = NO + N0 2 + 2H 2 S0 4

Covalent nitrosyl chloride N0C1 - orange-red gas (t pl= -65 ° C, t bale =-6 ° C), formed during chlorination of NO in the presence of activated carbon:

NO + C1 2 = 2N0C1

when nitrite interacts with hydrogen chloride:

NaN0 2 + 2HC1 = N0C1 + NaCl + H 2 0

or when replacing the anion in nitrosonium salts:

NaCl = N0C1 + NaHS0 4

Oxidizing properties are less characteristic of NO. For example, when interacting with strong reducing agents, nitrogen is formed:

2N0 + 2H 2 S = N 2 + 2S ↓ + 2H 2 0

On a rhodium catalyst, NO oxidizes carbon monoxide to carbon dioxide:

2N0 + 2CO = N 2 + 2C0 2

These catalysts are installed in the exhaust pipes of automobiles to avoid contamination of the atmosphere with toxic NO x gases.

When interacting with molten alkali, NO disproportionates:

6N0 + 4KON = N 2 + 4KN0 2 + 2H 2 0

Nitric oxide (III) N 2 0 3. This compound is very unstable and only exists at low temperatures. In the solid and liquid state (t pl = -100 ° C), this substance is colored bright blue; above O ° C it decomposes:

N 2 0 3 = N0 + N0 2

Unlike N 2 0 and NO, nitric oxide (II) is a typical acidic oxide; it dissolves in ice water to form a blue solution of nitrous acid;

N 2 0 3 + H 2 0 = 2HNO 2

When interacting with alkaline solutions, N 2 0 3 quantitatively transforms into nitrites:

N 2 0 3 + 2NaOH = 2NaN0 2 + H 2 0

In a strongly acidic medium, heterolytic decomposition of the NO-NO 2 bond occurs, as a result of which nitrosonium salts are formed:

N 2 0 3 + 3H 2 S0 4 = 2NO + + H 3 0 + + 3HSO 4

Upon cooling to -36 ° C, an equimolar mixture of NO and NO 2 oxides formed upon reduction with 50% HNO 3 with arsenic (III) oxide or starch condenses N 2 0 3:

2HN0 3 + As 2 0 3 + 2H 2 0 = 2H 3 As0 4 + N 2 0 3

1 / n (C 6 H 10 O 5) n + 12HN0 3 = 6C0 2 + 11H 2 0 + 6N 2 0 3

Nitrogen (IV) oxides: NO 2 and N 2 0 4. Nitrogen oxide (IV) in a wide temperature range exists in the form of an equilibrium mixture of monomer NO 2 and dimer N 2 0 4.

Equilibrium

2N0 2 ↔ N 2 0 4, ΔН =-57.2 kJ / mol

Brown gas Colorless gas

paramagnetic diamagnetic

highly dependent on temperature. Solid nitric oxide (IV) is colorless, since it consists exclusively of N 2 0 4 molecules. When it heats up to t, w= -12.8 ° C, a brown color appears, which increases with increasing temperature as the proportion of monomer in the mixture increases.

Nitric oxide (IV) (both monomer and dimer) is readily soluble in water and interacts with it. Since nitrogen compounds in even oxidation states do not exist in aqueous solutions, disproportionation into nitric and nitrous acids occurs:

N 2 0 4 + H 2 0 = HN0 3 + HN0 2

The latter is stable only in cold weather, and at room temperature and above it disproportionates to N0 and HN0 3, therefore, at room and higher temperatures, the reaction proceeds according to the equation

3N0 2 + H 2 0 = 2HN0 3 + NO

However, if a mixture of N0 2 and air is passed through the water, then only HN0 3 is formed:

2N0 2 + H 2 0 + 1/2 0 2 = 2HN0 3

Like NO, oxide N 2 0 4 is subject to one-electron oxidation with the formation of a cation nitroyl (nitronium) N0 2 , which has a linear structure and isoelectronic (16 e - for three atoms) C0 2. Nitroyl ion is also formed during the self-ionization of nitric acid:

2HN0 3 ↔ N0 2 + + NO 3 - + H 2 0

Dioxide NO 2 is a strong oxidizing agent, in the atmosphere of which carbon, sulfur, and many metals burn:

C + 2N0 2 = C0 2 + 2NO

In the gas phase, nitrogen dioxide oxidizes hydrogen chloride to chlorine:

2N0 2 + 4HC1 = 2NOC1 + 2H 2 0 + C1 2

N0 2 is obtained by the interaction of copper with hot concentrated nitric acid:

Cu + 4HN0 3 = Cu (N0 3) 2 + 2N0 2 + 2H 2 0

or by thermal decomposition (350-500 ° C) of carefully dried nitrates of heavy metals:

2Pb (N0 3) 2 → 2PbO + 4N0 2 + 0 2

The reaction is carried out in the presence of silicon dioxide, which binds the formed lead oxide into silicate PbSiO 3, thereby shifting the equilibrium to the right.

Nitric oxide (IV) is also formed during the oxidation of NO with oxygen:

2NO + 0 2 = 2N0 2, ΔН °= -114 kJ / mol

Interestingly, this reaction is reversible, and at 200 ° C the equilibrium is significantly shifted to the left.

Nitric oxide (V) N 2 0 5. Nitric anhydride N 2 0 5 is formed in the form of volatile (t subl = 32.3 ° C) colorless hygroscopic crystals when nitric acid vapor is passed through a column with phosphorus oxide (V):

4HN0 3 + Р 4 0 10 → 2N 2 0 5 + 4НР0 3

Solid N 2 0 5 is built of NO 2 + and N0 3 - ions, and in the gas phase and in solutions it consists of 0 2 N-O-N0 2 molecules. This substance is very unstable and decomposes within a few hours (half-life 10 hours), when heated - with an explosion:

2N 2 0 5 = 4N0 2 + 0 2

When N 2 0 5 is dissolved in water, nitric acid is formed.

Higher nitric oxide is a strong oxidizing agent, for example:

N 2 0 5 + I 2 = I 2 0 5 + N 2

In anhydrous acids (sulfuric, nitric, orthophosphoric, perchloric) N 2 0 5 decomposes, forming the nitronium cation N0 2:

N 2 0 5 + НСlO 4 = N0 2 + C10 4 - + HN0 3

Nitronium salts are strong oxidizing agents. When released into water, they hydrolyze:

N0 2 + C10 4 - + H 2 0 = HN0 3 + HC10 4

Nitroyl chloride N0 2 C1 (t pl = -145 ° C, t bale = -16 ° C) is a colorless gas formed when chlorine is passed over solid silver nitrate or when fuming nitric and chlorosulfonic acids react:

HN0 3 + ClSO 3 H = N0 2 C1 + H 2 S0 4

In an alkaline environment, it decomposes into hypochlorite and nitrite.

The content of the article

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, let us note that there is so much nitrogen in the atmosphere above each square kilometer of the earth's surface 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. Bound nitrogen is part of both organic and inorganic matter. The flora and fauna contain nitrogen bound to carbon and oxygen in proteins. In addition, nitrogen-containing inorganic compounds are known and can be obtained in large quantities, such as nitrates (NO 3 -), nitrites (NO 2 -), cyanides (CN -), nitrides (N 3–) and azides (N 3 - ).

Historical reference.

The experiments of A. Lavoisier, devoted to the study of the role of the atmosphere in maintaining life and combustion processes, confirmed the existence of a relatively inert substance in the atmosphere. Having failed to establish 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 nitrogen cycle.

The term "nitrogen fixation" means the process of fixing atmospheric nitrogen N 2. In nature, this can happen in two ways: either legumes, such as peas, clover and soybeans, accumulate nodules on their roots, in which the bacteria fixing nitrogen convert it to nitrates, or atmospheric nitrogen is oxidized by oxygen under conditions of a lightning discharge. S. Arrhenius found that up to 400 million tons of nitrogen is fixed in this way annually. In the atmosphere, nitrogen oxides combine with rainwater to form nitric and nitrous acids. In addition, it was found that with rain and snow, approx. 6700 g of 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 plants and convert them into animal proteins. After the death of animals and plants, their decomposition occurs, nitrogen compounds are converted 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. Thus, the nitrogen cycle occurs in nature, or the nitrogen cycle.

The structure of the nucleus and electron shells.

There are two stable nitrogen isotopes in nature: with a mass number of 14 (contains 7 protons and 7 neutrons) and with a mass number of 15 (contains 7 protons and 8 neutrons). Their ratio is 99.635: 0.365, so the atomic mass of nitrogen is 14.008. Unstable nitrogen isotopes 12 N, 13 N, 16 N, 17 N were obtained artificially. Schematically, the electronic structure of the nitrogen atom is as follows: 1 s 2 2s 2 2p x 1 2p y 1 2p z 1 . Consequently, there are 5 electrons on the outer (second) electron shell, which can participate in the formation of chemical bonds; nitrogen orbitals can also accept electrons, i.e. the formation of compounds with the oxidation state from (–III) to (V) is possible, and they are known.

Molecular nitrogen.

It was established from the definitions of the gas density that the nitrogen molecule is diatomic, i.e. the molecular formula of nitrogen is Nє N (or N 2). Two nitrogen atoms have three outer 2 p-electrons of each atom form a triple bond: N ::: N :, forming electron pairs. The measured interatomic N – N distance is 1.095 Å. As in the case of hydrogen ( cm... HYDROGEN), there are nitrogen molecules with different nuclear spins - symmetric and antisymmetric. At normal temperature, the ratio of symmetric and antisymmetric forms is 2: 1. Two modifications of nitrogen are known in the solid state: 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... tab. 1).

The dissociation energy of a mole (28.016 g or 6.023 × 10 23 molecules) of molecular nitrogen into atoms (N 2 2N) is approximately –225 kcal. Therefore, atomic nitrogen can be formed during a quiet electric discharge and is chemically more active than molecular nitrogen.

Receiving and applying.

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 permissible.

Nitrogen from the atmosphere.

Economically, the release of nitrogen from the atmosphere is due to the cheapness of the method of liquefying purified air (water vapor, CO 2, dust, and other impurities are removed). Successive cycles of compression, cooling and expansion of such air lead to its liquefaction. Liquid air is subjected to fractional distillation with a slow rise in temperature. Noble gases are released first, then nitrogen, and liquid oxygen remains. Purification is achieved by multiple fractionation processes. This method produces many millions of tons of nitrogen annually, mainly for the synthesis of ammonia, which is a feedstock in the technology for the production 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.

Small amounts of nitrogen can be obtained in the laboratory in various ways, by oxidizing ammonia or ammonium ion, for example:

The process of oxidation of the ammonium ion by the nitrite ion is very convenient:

Other methods are known - decomposition of azides upon heating, decomposition of ammonia with copper (II) oxide, interaction of nitrites with sulfamic acid or urea:

During the catalytic decomposition of ammonia at high temperatures, nitrogen can also be obtained:

Physical properties.

Some of the 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, ° С	–209,96
Boiling point, ° С	–195,8
Critical temperature, ° С	–147,1
Critical pressure, atm a	33,5
Critical density, g / cm 3 a	0,311
Specific heat, J / (molChK)	14.56 (15 ° C)
Pauling electronegativity	3
Covalent radius,	0,74
Crystalline radius,	1.4 (M 3–)
Ionization potential, V b
first	14,54
second	29,60
a Temperature and pressure at which the densities of liquid and gaseous nitrogen are the same. b The amount of energy required to remove the first external and the following electrons, per 1 mole of atomic nitrogen.

Chemical properties.

As already noted, the predominant property of nitrogen under normal conditions of temperature and pressure is its inertness, or low chemical activity. The electronic structure of nitrogen contains an electron pair for 2 s-level and three half filled 2 R-orbitals, so one nitrogen atom can bind no more than four other atoms, i.e. its coordination number is four. The small size of an atom also limits the number of atoms or groups of atoms that can be associated with it. Therefore, many compounds of other members of the VA subgroup either have no analogues among nitrogen compounds at all, or analogous nitrogen compounds are unstable. Thus, PCl 5 is a stable compound, while NCl 5 does not exist. A nitrogen atom is able to bind to 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 to form partially ionic nitrides M x N y... In these compounds, nitrogen is negatively charged. Table 2 shows the oxidation states and examples of the corresponding compounds.

Nitrides.

Nitrogen compounds with more electropositive elements, metals and non-metals - nitrides - are similar to carbides and hydrides. They can be divided, depending on the nature of the M – N bond, into ionic, covalent, and with an intermediate type of bond. As a rule, these are crystalline substances.

Ionic nitrides.

The bond in these compounds involves the transfer of electrons from metal to nitrogen with the formation of the N 3– ion. These nitrides include Li 3 N, Mg 3 N 2, Zn 3 N 2 and Cu 3 N 2. In addition to lithium, other alkali metals IA do not form the nitride subgroup. 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 (eg ammonia and hydrazine) are completely covalent, as are nitrogen halides (NF 3 and NCl 3). Covalent nitrides include, for example, Si 3 N 4, P 3 N 5 and BN - highly stable white substances, and BN has two allotropic modifications: hexagonal and diamond-like. The latter is formed at high pressures and temperatures and has a hardness close to that of diamond.

Nitrides with an intermediate type of bond.

Transition elements react with NH 3 at high temperatures to form an unusual class of compounds in which nitrogen atoms are distributed between regularly spaced metal atoms. There is no clear displacement of electrons in these compounds. Examples of such nitrides are Fe 4 N, W 2 N, Mo 2 N, Mn 3 N 2. These compounds are generally completely inert and have good electrical conductivity.

Hydrogen compounds of nitrogen.

Nitrogen and hydrogen interact to form compounds that vaguely resemble hydrocarbons. The stability of hydrogen nitrogen decreases with an increase in the number of nitrogen atoms in the chain, in contrast to hydrocarbons, which are also stable in long chains. The most important hydrogen nitrides are ammonia NH 3 and hydrazine N 2 H 4. They also include hydrazoic 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. USA produced approx. 13 million tons of ammonia annually (in terms of anhydrous ammonia).

Molecule structure.

The NH 3 molecule has an almost pyramidal structure. The H – N – H bond angle is 107 °, which is close to the tetrahedral angle of 109 °. An unshared electron pair is equivalent to an attached group; as a result, the coordination number of nitrogen is 4 and nitrogen is located in the center of the tetrahedron.

Ammonia properties.

Some of the physical properties of ammonia in comparison with water are given in table. 3.

The boiling and melting points of ammonia are much lower than those of water, despite the closeness of molecular weights and similarity of molecular structure. This is due to the relatively higher strength of intermolecular bonds in water than in ammonia (such an intermolecular bond is called hydrogen).

Ammonia as a solvent.

The high dielectric constant and dipole moment of liquid ammonia make it suitable for use as a solvent for polar or ionic inorganic substances. The ammonia solvent is intermediate 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 the solution according to the scheme

Blue is associated with the solvation and movement of electrons, or with the mobility of "holes" in a liquid. At a high sodium concentration in liquid ammonia, the solution takes on a bronze color and has a high electrical conductivity. Unbound alkali metal can be recovered 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 taking place in water:

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

Liquid ammonia as a solvent is advantageous in some cases when it is impossible to carry out reactions in water due to the rapid interaction of components with water (for example, oxidation and reduction). For example, in liquid ammonia, calcium reacts with KCl to form 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.

Getting ammonia.

Gaseous NH 3 is released from ammonium salts under the action of a strong base, for example, NaOH:

The method is applicable in laboratory conditions. Small-scale ammonia production is also based on the hydrolysis of nitrides, for example Mg 3 N 2, with water. Calcium cyanamide CaCN 2, when interacting with water, also forms ammonia. The main industrial method for producing ammonia is its catalytic synthesis from atmospheric nitrogen and hydrogen at high temperatures and pressures:

Hydrogen for this synthesis is obtained by thermal cracking of hydrocarbons, the action of water vapor on coal or iron, decomposition of alcohols with water vapor, or electrolysis of water. A lot of patents have been obtained for the synthesis of ammonia, differing in the conditions of the process (temperature, pressure, catalyst). There is a method of industrial production by thermal distillation of coal. The names of F. Gaber and K. Bosch are associated with the technological development of ammonia synthesis.

Table 4. COMPARISON OF REACTIONS IN AQUEOUS AND AMMONIA MEDIA
Water environment	Ammonia environment
Neutralization
OH - + H 3 O + ® 2H 2 O	NH 2 - + NH 4 + ® 2NH 3
Hydrolysis (protolysis)
PCl 5 + 3H 2 O POCl 3 + 2H 3 O + + 2Cl -	PCl 5 + 4NH 3 PNCl 2 + 3NH 4 + + 3Cl -
Substitution
Zn + 2H 3 O + ® Zn 2+ + 2H 2 O + H 2	Zn + 2NH 4 + ® Zn 2+ + 2NH 3 + H 2
Solvation (complexation)
Al 2 Cl 6 + 12H 2 O 2 3+ + 6Cl -	Al 2 Cl 6 + 12NH 3 2 3+ + 6Cl -
Amphotericity
Zn 2+ + 2OH - Zn (OH) 2	Zn 2+ + 2NH 2 - Zn (NH 2) 2
Zn (OH) 2 + 2H 3 O + Zn 2+ + 4H 2 O	Zn (NH 2) 2 + 2NH 4 + Zn 2+ + 4NH 3
Zn (OH) 2 + 2OH - Zn (OH) 4 2–	Zn (NH 2) 2 + 2NH 2 - Zn (NH 2) 4 2–

Chemical properties of ammonia.

In addition to the reactions mentioned in table. 4, ammonia reacts with water to form the compound NH 3 H H 2 O, which is often mistakenly thought to be ammonium hydroxide NH 4 OH; in fact, the existence of NH 4 OH in solution has not been proven. An aqueous solution of ammonia ("ammonia") consists mainly of NH 3, H 2 O and small concentrations of NH 4 + and OH - ions formed during dissociation

The main character of ammonia is explained by the presence of a lone electron pair of nitrogen: NH 3. Therefore, NH 3 is a Lewis base, which has a higher nucleophilic activity, manifested in the form of an association with a proton, or the nucleus of a hydrogen atom:

Any ion or molecule capable of accepting an electron pair (electrophilic compound) will react with NH 3 to form a coordination compound. For example:

Symbol M n+ represents a transition metal ion (B-subgroups of the periodic table, for example, Cu 2+, Mn 2+, etc.). 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, phosphate ammonium (NH 4) 3 PO 4. These salts are widely used in agriculture as fertilizers for introducing nitrogen into the soil. Ammonium nitrate is also used as an inexpensive explosive; it was first used with fuel oil (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 reacts with hydrides and nitrides to form amides:

Alkali metal amides (eg NaNH 2) react with N 2 O when heated to form azides:

Gaseous NH 3 reduces oxides of heavy metals to metals at high temperatures, apparently due to the hydrogen formed as a result of the decomposition of ammonia into N 2 and H 2:

The hydrogen atoms in the NH 3 molecule can be replaced by halogen. Iodine reacts with concentrated NH 3 solution to form a mixture of substances containing NI 3. This substance is very unstable and explodes at the slightest mechanical impact. When NH 3 reacts with Cl 2, chloramines NCl 3, NHCl 2 and NH 2 Cl are formed. When ammonia is exposed to sodium hypochlorite NaOCl (formed from NaOH and Cl 2), the final product is hydrazine:

Hydrazine.

The above reactions represent a method for preparing hydrazine monohydrate with the composition N 2 H 4 CH H 2 O. Anhydrous hydrazine is formed by special distillation of the monohydrate with BaO or other dehydrating substances. In terms of properties, hydrazine slightly resembles hydrogen peroxide H 2 O 2. Pure anhydrous hydrazine is a colorless hygroscopic liquid boiling at 113.5 ° C; well soluble in water, forming a weak base

In an acidic medium (H +) hydrazine forms soluble hydrazonium salts of the + X - type. The ease with which hydrazine and some of its derivatives (for example, methylhydrazine) react with oxygen allows it to be used as a component of liquid propellant. Hydrazine and all of its derivatives are highly toxic.

Nitrogen oxides.

In compounds with oxygen, nitrogen exhibits all oxidation states, forming oxides: N 2 O, NO, N 2 O 3, NO 2 (N 2 O 4), N 2 O 5. There is scant information on the formation of nitrogen peroxides (NO 3, NO 4). 2HNO 2. Pure N 2 O 3 can be obtained as a blue liquid at low temperatures (-20

At room temperature, NO 2 is a dark brown gas that has magnetic properties due to the presence of an unpaired electron. At temperatures below 0 ° C, the NO 2 molecule dimerizes to dinitrogen tetroxide, and at –9.3 ° C the dimerization proceeds completely: 2NO 2 N 2 O 4. In the liquid state, only 1% NO 2 is undimerized, and at 100 ° C it remains in the form of a dimer of 10% N 2 O 4.

NO 2 (or N 2 O 4) reacts in warm water to form nitric acid: 3NO 2 + H 2 O = 2HNO 3 + NO. The NO 2 technology is therefore very important as an intermediate stage in the production of an industrially important product - nitric acid.

Nitric oxide (V)

N 2 O 5 ( outdated... nitric acid anhydride) is a white crystalline substance, obtained by dehydration of nitric acid in the presence of phosphorus oxide P 4 O 10:

2MX + H 2 N 2 O 2. Evaporation of the solution produces a white explosive with the assumed structure H – O – N = N – O – H.

Nitrous acid

HNO 2 does not exist in its pure form, however, aqueous solutions of its low concentration are formed when sulfuric acid is added to barium nitrite:

Nitrous acid is also formed by dissolving an equimolar mixture of NO and NO 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 is H – O – N = O); it can be both an oxidizing agent and a reducing agent. Under the action of reducing agents, it is usually reduced to NO, and when interacting with oxidants, it is oxidized to nitric acid.

The dissolution rate 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 readily soluble in water, except for silver nitrite. NaNO 2 is used in the production of dyes.

Nitric acid

HNO 3 is one of the most important inorganic products of the main chemical industry. It is used in the technologies of many other inorganic and organic substances, for example, explosives, fertilizers, polymers and fibers, dyes, pharmaceuticals, etc.

Literature:

Azotchik's Handbook... M., 1969
B.V. Nekrasov Fundamentals of General Chemistry... M., 1973
Nitrogen fixation problems. Inorganic and physical chemistry... M., 1982



Introduction

If you take a close look at nitrogen in the periodic table of chemical elements of D.I.Mendeleev, you will notice that it has a variable valence. This means that nitrogen forms several binary compounds with oxygen at once. Some of them have been discovered recently, and some have been studied far and wide. There are unstable and stable nitrogen oxides. The chemical properties of each of these substances are completely different, therefore, when studying them, at least five nitrogen oxides must be considered. This is what we will talk about in today's article.

Nitric oxide (I)

The formula is N 2 O. It can sometimes be called nitrogen oxonitride, dinitrogen oxide, nitrous oxide or laughing gas.

Properties

Under normal conditions, it is represented by a colorless gas with a sweetish odor. It can be dissolved by water, ethanol, ether and sulfuric acid. If gaseous monovalent nitrogen oxide is heated to room temperature under a pressure of 40 atmospheres, then it thickens to a colorless liquid. It is a non-salt-forming oxide that decomposes during heating and shows itself in reactions as a reducing agent.

Receiving

This oxide is formed when heated dry. Another way to obtain it is by thermal decomposition of a mixture of "sulfamic + nitric acid".

Application

Used as an inhalation anesthetic, the food industry knows this oxide as an additive E942. It also improves the technical characteristics of internal combustion engines.

Nitric oxide (II)

Formula - NO. Occurs under the names of nitrogen monoxide, nitric oxide and nitrosyl radical

Properties

Under normal conditions, it appears as a colorless gas that is poorly soluble in water. It is difficult to liquefy it, but it is blue in solid and liquid states. This oxide can be oxidized by atmospheric oxygen.

Receiving

It is quite simple to obtain it, for this you need to heat a mixture of nitrogen and oxygen to 1200-1300 ° C. In laboratory conditions, it is formed at once in several experiments:

• Reaction of copper and 30% nitric acid solution.
• Reaction between sodium nitrite and hydrochloric acid.
• Reaction of nitrous and hydroiodic acids.

Application

This is one of the substances from which nitric acid is obtained.

Nitric oxide (III)

Formula - N 2 O 3. It can also be called nitrous anhydride and nitrogen sesquioxide.

Properties

Under normal conditions it is a liquid that is blue, and under standard conditions it is a colorless gas. Pure oxide exists only in a solid state of aggregation.

Receiving

Formed by the interaction of 50% nitric acid and solid oxide of trivalent arsenic (it can also be replaced with starch).

Application

With the help of this substance, its salts are obtained in laboratories.

Nitric oxide (IV)

Formula - NO 2. It can also be called nitrogen dioxide or brown gas.

Properties

The last name corresponds to one of its properties. After all, this oxide has the form of either a red-brown gas or a yellowish liquid. It is characterized by high chemical activity.

Receiving

This oxide is obtained by the interaction of nitric acid and copper, as well as during the thermal decomposition of lead nitrate.

Application

With the help of it, sulfuric and nitric acids are produced, liquid and mixed

Nitric oxide (V)

Formula - N 2 O 5. May be referred to as dinitrogen pentoxide, nitroyl nitrate, or nitric anhydride.

Properties

Has the appearance of colorless and highly volatile crystals. They can melt at a temperature of 32.3 ° C.

Receiving

This oxide is formed by several reactions:

• Dehydration of nitric acid with pentavalent phosphorus oxide.
• Passage of dry chlorine over
• Interaction of ozone with tetravalent nitrogen oxide.

Application

Due to its extreme instability, it is not used anywhere in its pure form.

Conclusion

There are nine nitrogen oxides in chemistry, the above are only the classic compounds of this element. The other four are, as already mentioned, unstable substances. However, they are all united by one property - high toxicity. Emissions of nitrogen oxides into the atmosphere lead to a deterioration in the health of people living in the vicinity of industrial chemical plants. Symptoms of poisoning with any of these substances are toxic pulmonary edema, disruption of the central nervous system and blood damage, the cause of which is the binding of hemoglobin. Therefore, nitrogen oxides must be handled with care and, in most cases, protective equipment must be used.

 

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