Compounds with non-metals

All nitrogen halides NG 3 are known. Trifluoride NF 3 is obtained by reacting fluorine with ammonia:

3F 2 + 4NH 3 = 3 NH 4 F + NF 3

Nitrogen trifluoride is a colorless toxic gas whose molecules have a pyramidal structure. Fluorine atoms are located at the base of the pyramid, and the top is occupied by a nitrogen atom with a lone electron pair. NF 3 is very resistant to various chemicals and heat.

The remaining nitrogen trihalides are endothermic and therefore unstable and reactive. NCl 3 is formed by passing chlorine gas into a strong solution of ammonium chloride:

3Cl 2 + NH 4 Cl = 4HCl + NCl 3

Nitrogen trichloride is a highly volatile (t boiling point = 71 degrees C) liquid with a pungent odor. A slight heating or impact is accompanied by an explosion releasing a large amount of heat. In this case, NCl 3 breaks down into elements. Trihalides NBr 3 and NI 3 are even less stable.

Nitrogen derivatives with chalcogens are very unstable due to their strong endothermicity. All of them are poorly studied and explode when heated and impacted.

Connections to metals

Salt-like nitrides are obtained by direct synthesis from metals and nitrogen. Salt-like nitrides decompose with water and dilute acids:

Mg 3 N 2 + 6N 2 = 3 Mg(OH) 2 + 2NH 3

Ca 3 N 2 + 8HCl = 3CaCl 2 + 2NH 4 Cl

Both reactions prove the basic nature of active metal nitrides.

Metal-like nitrides are produced by heating metals in an atmosphere of nitrogen or ammonia. Oxides, halides and hydrides of transition metals can be used as starting materials:

2Ta + N 2 = 2TaN; Mn 2 O 3 + 2NH 3 = 2 MnN + 3H 2 O

CrCl 3 + NH 3 = CrN + 3HCl; 2TiN 2 + 2NH 3 = 2TiN +5H 2

Application of nitrogen and nitrogen-containing compounds

The scope of application of nitrogen is very wide - the production of fertilizers, explosives, ammonia, which is used in medicine. Nitrogen-containing fertilizers are the most valuable. Such fertilizers include ammonium nitrate, urea, ammonia, and sodium nitrate. Nitrogen is an integral part of protein molecules, which is why plants need it for normal growth and development. Such an important compound of nitrogen with hydrogen as ammonia is used in refrigeration units; ammonia, circulating through a closed pipe system, takes away a large amount of heat during its evaporation. Potassium nitrate is used to produce black powder, and gunpowder is used in hunting rifles and for exploration of ore deposits lying underground. Black powder is obtained from pyroxylin, an ester of cellulose and nitric acid. Organic explosives based on nitrogen are used to build tunnels in the mountains (TNT, nitroglycerin).

Being in nature.

Nitrogen occurs in nature mainly in a free state. In air, its volume fraction is 78.09%, and its mass fraction is 75.6%. Nitrogen compounds are found in small quantities in soils. Nitrogen is part of proteins and many natural organic compounds. The total nitrogen content in the earth's crust is 0.01%.

Receipt.

In technology, nitrogen is obtained from liquid air. As you know, air is a mixture of gases, mainly nitrogen and oxygen. Dry air at the Earth's surface contains (in volume fractions): nitrogen 78.09%, oxygen 20.95%, noble gases 0.93%, carbon monoxide (IV) 0.03%, as well as random impurities - dust, microorganisms , hydrogen sulfide, sulfur oxide (IV), etc. To obtain nitrogen, air is transferred to a liquid state, and then nitrogen is separated from less volatile oxygen by evaporation (i.e. boiling point of nitrogen -195.8 °C, oxygen -183 °C). The nitrogen obtained in this way contains impurities of noble gases (mainly argon). Pure nitrogen can be obtained in laboratory conditions by decomposing ammonium nitrite when heated:

NH 4 NO 2 = N 2 + 2H 2 O

Physical properties. Nitrogen is a colorless, odorless and tasteless gas, lighter than air. Solubility in water is less than that of oxygen: at 20 0 C, 15.4 ml of nitrogen dissolves in 1 liter of water (oxygen 31 ml). Therefore, in air dissolved in water, the oxygen content relative to nitrogen is greater than in the atmosphere. The low solubility of nitrogen in water, as well as its very low boiling point, are explained by very weak intermolecular interactions both between nitrogen and water molecules, and between nitrogen molecules.

Natural nitrogen consists of two stable isotopes with mass numbers 14 (99.64%) and 15 (0.36%).

Chemical properties.

At room temperature, nitrogen combines directly only with lithium:

6Li + N 2 = 2Li 3 N

It reacts with other metals only at high temperatures, forming nitrides. For example:

3Ca + N 2 = Ca 3 N 2, 2Al + N 2 = 2AlN

Nitrogen combines with hydrogen in the presence of a catalyst at high blood pressure and temperature:

N2 + 3H2 = 2NH3

At the temperature of the electric arc (3000-4000 degrees), nitrogen combines with oxygen:

Application. Nitrogen is used in large quantities to produce ammonia. Widely used to create an inert environment - filling incandescent electric lamps and free space in mercury thermometers, when pumping flammable liquids. It is used to nitrate the surface of steel products, i.e. saturate their surface with nitrogen at high temperatures. As a result, iron nitrides are formed in the surface layer, which impart greater hardness to the steel. This steel can withstand heating up to 500 °C without losing its hardness.

Nitrogen is important for the life of plants and animals, since it is part of protein substances. Nitrogen compounds are used in the production of mineral fertilizers, explosives and in many industries.

Question No. 48.

Ammonia, its properties, methods of production. The use of ammonia in the national economy. Ammonium hydroxide. Ammonium salts, their properties and applications. Nitrogen fertilizers with the ammonium form of nitrogen. Qualitative reaction to ammonium ion.

Ammonia – a colorless gas with a characteristic odor, almost twice as light as air. When pressure is increased or cooled, it easily liquefies into a colorless liquid. Ammonia is very soluble in water. A solution of ammonia in water is called ammonia water or ammonia. When boiling, dissolved ammonia evaporates from the solution.

Chemical properties.

Interaction with acids:

NH 3 + HCl = NH 4 Cl, NH 3 + H 3 PO 4 = NH 4 H 2 PO 4

Interaction with oxygen:

4NH 3 + 3O 2 = 2N 2 + 6H 2 O

Copper recovery:

3CuO + 2NH 3 = 3Cu + N 2 + 3H 2 O

Receipt.

2NH 4 Cl + Ca(OH) 2 = CaCl 2 + 2NH 3 + 2H 2 O

N2 + 3H2 = 2NH3

Application.

Liquid ammonia and its aqueous solutions are used as liquid fertilizer.

Ammonium hydroxide (ammonium hydroxide) – NH 4 OH

Ammonium salts and their properties. Ammonium salts consist of an ammonium cation and an acid anion. They are similar in structure to the corresponding salts of singly charged metal ions. Ammonium salts are obtained by reacting ammonia or its aqueous solutions with acids. For example:

NH 3 + HNO 3 = NH 4 NO 3

They exhibit the general properties of salts, i.e. interact with solutions of alkalis, acids and other salts:

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

2NH 4 Cl + H 2 SO 4 = (NH 4) 2 SO 4 + 2HCl

(NH 4) 2 SO 4 + BaCl 2 = BaSO 4 + 2NH 4 Cl

Application. Ammonium nitrate (ammonium nitrate) NH4NO3 is used as a nitrogen fertilizer and for the production of explosives - ammonites;

Ammonium sulfate (NH4)2SO4 - as a cheap nitrogen fertilizer;

Ammonium bicarbonate NH4HCO3 and ammonium carbonate (NH4)2CO3 - in the food industry in the production of flour confectionery products as a chemical leavening agent, in dyeing fabrics, in the production of vitamins, in medicine;

Ammonium chloride (ammonia) NH4Cl - in galvanic cells (dry batteries), during soldering and tinning, in the textile industry, as a fertilizer, in veterinary medicine.

Ammonium (ammonia) fertilizers contain nitrogen in the form of ammonium ion and have an acidifying effect on the soil, which leads to a deterioration in its properties and to less effective fertilizers, especially when applied regularly on unlimed, infertile soils. But these fertilizers also have their advantages: ammonium is much less susceptible to leaching, since it is fixed by soil particles and absorbed by microorganisms, and, in addition, the process of nitrophification occurs with it in the soil, i.e. conversion by microorganisms into nitrates. Of the ammonium fertilizers, ammonium chloride is the least suitable for vegetable crops as it contains quite a lot of chlorine.

Qualitative reaction to ammonium ion.

A very important property of ammonium salts is their interaction with alkali solutions. This reaction is detected by ammonium salts (ammonium ion) by the smell of ammonia released or by the appearance of a blue color on wet red litmus paper:

NH 4 + + OH - = NH 3 + H 2 O

The chemical element nitrogen forms only one simple substance. This substance is gaseous and is formed by diatomic molecules, i.e. has the formula N 2. Despite the fact that the chemical element nitrogen has high electronegativity, molecular nitrogen N2 is an extremely inert substance. This fact is due to the fact that the nitrogen molecule contains an extremely strong triple bond (N≡N). For this reason, almost all reactions with nitrogen occur only at elevated temperatures.

Interaction of nitrogen with metals

The only substance that reacts with nitrogen under normal conditions is lithium:

An interesting fact is that with the rest of the active metals, i.e. alkaline and alkaline earth, nitrogen reacts only when heated:

The interaction of nitrogen with metals of medium and low activity (except Pt and Au) is also possible, but requires incomparably higher temperatures.

Nitrides of active metals are easily hydrolyzed by water:

As well as acid solutions, for example:

Interaction of nitrogen with nonmetals

Nitrogen reacts with hydrogen when heated in the presence of catalysts. The reaction is reversible, therefore, to increase the yield of ammonia in industry, the process is carried out at high pressure:

As a reducing agent, nitrogen reacts with fluorine and oxygen. The reaction with fluorine occurs under the action of an electric discharge:

The reaction with oxygen occurs under the influence of an electric discharge or at a temperature of more than 2000 o C and is reversible:

Of the non-metals, nitrogen does not react with halogens and sulfur.

Interaction of nitrogen with complex substances

Chemical properties of phosphorus

There are several allotropic modifications of phosphorus, in particular white phosphorus, red phosphorus and black phosphorus.

White phosphorus is formed by tetraatomic P4 molecules and is not a stable modification of phosphorus. Poisonous. At room temperature soft and, like wax, easy to cut with a knife. It oxidizes slowly in air, and due to the peculiarities of the mechanism of such oxidation, it glows in the dark (the phenomenon of chemiluminescence). Even with low heating, spontaneous ignition of white phosphorus is possible.

Of all allotropic modifications, white phosphorus is the most active.

Red phosphorus consists of long molecules of variable composition Pn. Some sources indicate that it has an atomic structure, but it is more correct to consider its structure molecular. Due to its structural features, it is a less active substance compared to white phosphorus; in particular, unlike white phosphorus, it oxidizes much more slowly in air and requires ignition to ignite.

Black phosphorus consists of continuous chains of P n and has a layered structure similar to the structure of graphite, which is why it looks similar to it. This allotropic modification has an atomic structure. The most stable of all allotropic modifications of phosphorus, the most chemically passive. For this reason, discussed below Chemical properties Phosphorus should be classified primarily as white and red phosphorus.

Interaction of phosphorus with non-metals

The reactivity of phosphorus is higher than that of nitrogen. Thus, phosphorus is able to burn after ignition under normal conditions, forming acidic oxide P 2 O 5:

and with a lack of oxygen, phosphorus (III) oxide:

The reaction with halogens is also intense. Thus, during the chlorination and bromination of phosphorus, depending on the proportions of the reagents, phosphorus trihalides or pentahalides are formed:

Due to the significantly weaker oxidizing properties of iodine compared to other halogens, oxidation of phosphorus with iodine is possible only to the oxidation state +3:

Unlike nitrogen phosphorus does not react with hydrogen.

Interaction of phosphorus with metals

Phosphorus reacts when heated with active metals and metals of intermediate activity to form phosphides:

Phosphides of active metals, like nitrides, are hydrolyzed by water:

As well as aqueous solutions of non-oxidizing acids:

Interaction of phosphorus with complex substances

Phosphorus is oxidized by oxidizing acids, in particular concentrated nitric and sulfuric acids:

You should know that white phosphorus reacts with aqueous solutions of alkalis. However, due to the specificity, the ability to write equations for such interactions on the Unified State Exam in chemistry has not yet been required.

However, for those who claim 100 points, for their own peace of mind, you can remember the following features of the interaction of phosphorus with alkali solutions in the cold and when heated.

In the cold, the interaction of white phosphorus with alkali solutions proceeds slowly. The reaction is accompanied by the formation of a gas with the smell of rotten fish - phosphine and a compound with a rare oxidation state of phosphorus +1:

When white phosphorus reacts with a concentrated alkali solution during boiling, hydrogen is released and phosphite is formed:

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

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

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

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

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

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

N 2

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

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

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

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

N 2 + 3F 2 = 2NF 3, N 2 + O 2 ↔ 2NO

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

N 2 + 3H 2 ↔ 2NH 3 + 92 kJ

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

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

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

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

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

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

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

N.H. 3

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

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

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

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

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

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

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

Nitrogen oxides

Nitrogen monoxideNO

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

Nitrogen dioxideNO 2

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

Dianitrogen oxideN 2 O

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

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

Dianitrogen trioxideN 2 O 3

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

Dianitrogen pentoxide N 2 O 5

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

Nitrites and nitrates

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

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

2KNO 3 = 2KNO 2 + O 2 (400-500 ˚C)

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

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

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

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

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

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

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

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