Compounds with non-metals

All nitrogen halides NГ 3 are known. Trifluoride NF 3 is obtained by the interaction of fluorine with ammonia:

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

Nitrogen trifluoride is a colorless toxic gas, the molecules of which have a pyramidal structure. At the base of the pyramid are dislocated fluorine atoms, and the top is occupied by a nitrogen atom with a lone electron pair. NF 3 is very resistant to various chemicals and to heating.

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

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

Nitrogen trichloride is a highly volatile (t bale = 71 deg. C) liquid with a pungent odor. A slight heating or shock is accompanied by an explosion with the release of a large amount of heat. In this case, NCl 3 decomposes into elements. Trihalides NBr 3 and NI 3 are even less stable.

Nitrogen derivatives with chalcogenes are very unstable due to their strong endothermicity. All of them are poorly studied; they explode when heated and hit.

Compounds with metals

Salt 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 = 3Mg (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 obtained by heating metals in a nitrogen or ammonia atmosphere. As starting materials, oxides, halides and hydrides of transition metals can be used:

2Ta + N 2 = 2TaN; Мn 2 О 3 + 2NH 3 = 2МnN + 3Н 2 О

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

The use of nitrogen and nitrogen-containing compounds

The area of ​​application of nitrogen is very large - the production of fertilizers, explosives, ammonia, which is used in medicine. Fertilizers containing nitrogen are the most valuable. Such fertilizers include ammonium nitrate, urea, ammonia, 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 system of pipes, during its evaporation takes up a large amount of heat. Potash nitrate is used for the production of black powder, and gunpowder is used in hunting rifles, for prospecting for ore minerals that lie underground. Black powder is obtained from pyroxylin - an ester of cellulose and nitric acid. Nitrogen-based organic explosives are used for tunneling in the mountains (TNT, nitroglycerin).

The chemical element nitrogen forms only one simple substance. This substance is gaseous and formed by diatomic molecules, i.e. has the formula N 2. Despite the fact that the chemical element nitrogen has a high electronegativity, molecular nitrogen N 2 is an extremely inert substance. This fact is due to the fact that an extremely strong triple bond (N≡N) takes place in the nitrogen molecule. For this reason, practically all reactions with nitrogen proceed 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 for Pt and Au) is also possible, however, it requires incomparably higher temperatures.

Nitrides of active metals are easily hydrolyzed with water:

And also with acid solutions, for example:

Interaction of nitrogen with non-metals

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. With fluorine, the reaction proceeds under the action of an electric discharge:

With oxygen, the reaction proceeds under the action of an electric discharge or at a temperature of more than 2000 ° C and is reversible:

Of 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, notably white phosphorus, red phosphorus, and black phosphorus.

White phosphorus is formed by tetraatomic P 4 molecules and is not a stable modification of phosphorus. Poisonous. At room temperature, it is soft and like a wax, it can be easily cut with a knife. It slowly oxidizes 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 the allotropic modifications, white phosphorus is the most active.

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

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

Interaction of phosphorus with non-metals

The reactivity of phosphorus is higher than that of nitrogen. So, phosphorus is capable of burning after ignition under normal conditions, forming an acidic oxide Р 2 O 5:

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

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

Due to the significantly weaker oxidizing properties of iodine in comparison with other halogens, it is possible for phosphorus to be oxidized with iodine only up 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 medium activity to form phosphides:

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

And also with 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 be aware that white phosphorus reacts with aqueous solutions of alkalis. However, due to the specificity, the ability to write down the equations of such interactions on the exam in chemistry has not yet been required.

Nevertheless, 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 interacts with a concentrated alkali solution, hydrogen is released during boiling and phosphite is formed:

Being in nature.

Nitrogen is found 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 amounts in soils. Nitrogen is a component of protein substances and many natural organic compounds. The total nitrogen content in the earth's crust is 0.01%.

Receiving.

In technology, nitrogen is obtained from liquid air. As you know, air is a mixture of gases, mainly nitrogen and oxygen. Dry air near 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 accidental impurities -, dust, microorganisms , hydrogen sulfide, sulfur (IV) oxide, etc. To obtain nitrogen, the air is converted into a liquid state, and then the nitrogen is separated from less volatile oxygen by evaporation (bp. nitrogen -195.8 ° C, oxygen -183 ° C). The nitrogen obtained in this way contains admixtures of noble gases (mainly argon). Pure nitrogen can be obtained under 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 ° C, 15.4 ml of nitrogen (oxygen 31 ml) dissolves in 1 liter of water. Therefore, in the air dissolved in water, the oxygen content in relation 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 only binds directly to 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 pressure and temperature:

N 2 + 3H 2 = 2NH 3

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

Application. Nitrogen is used in large quantities to obtain ammonia. It is widely used to create an inert environment - filling electric incandescent lamps and free space in mercury thermometers when pumping flammable liquids. The surface of steel products is nitrided with it, i.e. saturate their surface with nitrogen at a high temperature. As a result, iron nitrides are formed in the surface layer, which impart greater hardness to the steel. Such 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 number 48.

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

Ammonia - a colorless gas with a characteristic odor, almost two times lighter than air. With increasing pressure or cooling, 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 boiled, 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

Receiving.

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

N 2 + 3H 2 = 2NH 3

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 are composed of an ammonium cation and an acid anion. In structure, they are similar to the corresponding salts of singly charged metal ions. Ammonium salts are obtained by the interaction of 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 nitrogen fertilizer and for the manufacture 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 as a chemical baking powder, in fabric dyeing, in the production of vitamins, in medicine;

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

Ammonium (ammonia) fertilizers contain nitrogen in the form of an ammonium ion and have an acidifying effect on the soil, which leads to a deterioration in its properties and to a lower efficiency of fertilizers, especially with regular application on unfrozen low-fertile 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 nitrophication occurs with it in the soil, i.e. transformation 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 released ammonia or by the appearance of a blue coloration of a wet red litmus paper:

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

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

Nitrogen oxidation scale:
+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), the third after F and O. It exhibits typical non-metallic (acidic) properties, while forming various oxygen-containing acids, salts and binary compounds, as well as ammonium cation NH 4 and its salts.

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

N 2

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

A colorless, odorless and tasteless gas that condenses into a colorless liquid (unlike O 2).

The main constituent of air is 78.09% by volume, 75.52% by mass. Nitrogen boils off from liquid air earlier than oxygen. It is 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, N 2 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 for producing ammonia takes place at a temperature of 200˚C, under a pressure of up to 350 atm and always in the presence of a catalyst (Fe, F 2 O 3, FeO, in the laboratory at Pt)

N 2 + 3H 2 ↔ 2NH 3 + 92 kJ

In accordance with the Le Chatelier principle, the increase in ammonia yield should occur with increasing pressure and decreasing temperature. However, the reaction rate at low temperatures is very low, therefore the process is carried out at 450-500 ˚C, reaching a 15% yield of ammonia. Unreacted N 2 and H 2 are recycled to the reactor and thereby increase the rate of reaction.

Nitrogen is chemically passive towards acids and alkalis and does not support combustion.

Receiving v industry- fractional distillation of liquid air or removal of oxygen from the air by chemical means, for example, by the reaction 2C (coke) + O 2 = 2CO when heated. In these cases, nitrogen is obtained, containing also admixtures of noble gases (mainly argon).

In the laboratory, small amounts of chemically pure nitrogen can be obtained by the contamination 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)

It is used for the synthesis of ammonia. Nitric acid and other nitrogen-containing products as an inert medium for chemical and metallurgical processes and storage of flammable substances.

NH 3

A binary compound, the oxidation state of nitrogen is - 3. Colorless gas with a pungent 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 in the NH 3 molecule in nitrogen in the sp 3 -hybrid orbital determines the characteristic reaction of the addition of a hydrogen cation, with the formation of a cation ammonium NH 4. It liquefies under excess pressure at room temperature. In the liquid state, it is associated due to hydrogen bonds. Thermally unstable. Let's well dissolve in water (more than 700 l / 1 l H 2 O at 20˚C); the proportion in a saturated solution is 34% by weight and 99% by volume, pH = 11.8.

Highly reactive, prone to addition reactions. Burns in oxygen, reacts with acids. Shows reducing (due to N -3) and oxidizing (due to H +1) properties. Dried only with calcium oxide.

Qualitative reactions - the formation of white "smoke" in 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. It is 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)
Receiving. V 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 followed by drying the gas.
In industry ammonia is obtained from nitrogen with hydrogen. Produced by the industry either in liquefied form or in the form of a concentrated aqueous solution under the technical name ammonia water.

Ammonia hydrateNH 3 * H 2 O. Intermolecular compound. White, in the crystal lattice there are NH 3 and H 2 O molecules bound by a weak hydrogen bond. Present in aqueous ammonia solution, 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 with strong acids. Shows reducing properties (due to N -3) in a concentrated solution. It enters into the reaction of ion exchange and complexation.

Qualitative reaction- formation of white "smoke" on contact with gaseous HCl. It is used to create a slightly alkaline medium 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% of NH 4 OH ions (due to the dissociation of the hydrate); 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 (dil.) = 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
Diluted 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. A radical, contains a covalent σπ-bond (N꞊O), in the solid state is an N 2 O 2 dimer with an N-N bond. Extremely thermally stable. Sensitive to oxygen in the air (turns brown). It is slightly soluble in water and does not react with it. Chemically passive towards acids and alkalis. Reacts with metals and non-metals when heated. highly reactive mixture of NO and NO 2 ("nitrous gases"). An intermediate in the synthesis of nitric acid.
Equations of the most important reactions:
2NO + O 2 (gas) = ​​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)
Receiving 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 reduction of nitrates:
2NaNO 2 + 2H 2 SO 4 + 2NaI = 2 NO + I 2 ↓ + 2 H 2 O + 2Na 2 SO 4

Nitrogen dioxideNO 2

Acidic oxide, conventionally corresponds to two acids - HNO 2 and HNO 3 (acid for N 4 does not exist). Brown gas, a monomer of NO 2 at room temperature, in the cold, a liquid colorless dimer of N 2 O 4 (dinitrogen tetroxide). Reacts completely with water, alkalis. Very strong oxidizing agent, corrosive to metals. It is used for the synthesis of nitric acid and anhydrous nitrates, as an oxidizing agent for rocket fuel, 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 (dil.) = 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)
Receiving: v industry - oxidation of NO with atmospheric oxygen, in laboratories- interaction of concentrated nitric acid with reducing agents:
6HNO 3 (conc., Horizontal) + S = H 2 SO 4 + 6NO 2 + 2H 2 O
5HNO 3 (conc., Horizontal) + P (red) = H 3 PO 4 + 5NO 2 + H 2 O
2HNO 3 (conc., Hot.) + SO 2 = H 2 SO 4 + 2 NO 2

Dinitrogen oxideN 2 O

A colorless gas with a pleasant odor ("laughing gas"), N꞊N꞊O, the formal oxidation state of nitrogen is +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)
Received 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.

Dinitrogen trioxideN 2 O 3

At low temperatures, blue liquid, ON꞊NO 2, formal nitrogen oxidation state +3. At 20 ˚C, it decomposes by 90% into a mixture of colorless NO and brown NO 2 ("nitrous gases", industrial smoke - "fox's 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 gives HNO 2 salts, for example NaNO 2.
Obtained by interaction of 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" and environmentally hazardous, act as catalysts for the destruction of the ozone layer of the atmosphere.

Dinitrogen pentoxide N 2 O 5

Colorless, solid, O 2 N - O - NO 2, the oxidation state of nitrogen is +5. At room temperature, it decomposes into NO 2 and O 2 in 10 hours. Reacts with water and alkalis as acidic oxide:
N 2 O 5 + H 2 O = 2HNO 3
N 2 O 5 + 2NaOH = 2NaNO 3 + H 2
Received by dehydration of fuming nitric acid:
2HNO 3 + P 2 O 5 = N 2 O 5 + 2HPO 3
or oxidation of NO 2 with ozone at -78˚C:
2NO 2 + O 3 = N 2 O 5 + O 2

Nitrite and nitrate

Potassium nitriteKNO 2 ... White, hygroscopic. Melts without decomposition. Resistant to dry air. Let's very well dissolve in water (forming a colorless solution), hydrolyzed by anion. Typical oxidizing and reducing agent in acidic environment, reacts very slowly in alkaline environment. It enters into ion exchange reactions. Qualitative reactions for NO 2 ion - discoloration of a violet MnO 4 solution and the appearance of a black precipitate when I ions are added. It is used in the production of dyes, as an analytical reagent for amino acids and iodides, a component of photographic reagents.
equation of the most important reactions:
2KNO 2 (s) + 2HNO 3 (conc.) = NO 2 + NO + H 2 O + 2KNO 3
2KNO 2 (dil.) + O 2 (gas) → 2KNO 3 (60-80 ˚C)
KNO 2 + H 2 O + Br 2 = KNO 3 + 2HBr
5NO 2 - + 6H + + 2MnO 4 - (phiol.) = 5NO 3 - + 2Mn 2+ (bts.) + 3H 2 O
3 NO 2 - + 8H + + CrO 7 2- = 3NO 3 - + 2Cr 3+ + 4H 2 O
NO 2 - (sat.) + NH 4 + (sat.) = N 2 + 2H 2 O
2NO 2 - + 4H + + 2I - (bts.) = 2NO + I 2 (black) ↓ = 2H 2 O
NO 2 - (expanded) + Ag + = AgNO 2 (light yellow) ↓
Receiving vindustry- recovery of potassium nitrate in the processes:
KNO 3 + Pb = KNO 2+ PbO (350-400˚C)
KNO 3 (conc.) + Pb (sponge) + H 2 O = KNO 2+ Pb (OH) 2 ↓
3 KNO 3 + CaO + SO 2 = 2 KNO 2+ CaSO 4 (300 ˚C)

H itrat potassium KNO 3
Technical name potash, or indian salt , saltpeter. White, melts without decomposition upon further heating decomposes. Resistant to air. Let's well dissolve in water (with high endo-effect, = -36 kJ), no hydrolysis. Strong oxidizing agent during fusion (due to the release of atomic oxygen). In solution, it is reduced only with atomic hydrogen (in an acidic medium to KNO 2, in an alkaline medium to NH 3). It is used in glass production as a food preservative, 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)

Receiving: in industry
4KOH (hot) + 4NO 2 + O 2 = 4KNO 3 + 2H 2 O

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