Finding in nature.

Nitrogen occurs in nature mainly in the 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 constituent of proteins and many natural organic compounds. The total nitrogen content in the earth's crust is 0.01%.

Receipt.

In engineering, 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 by evaporation from less volatile oxygen (bp. 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 the laboratory by decomposing ammonium nitrite when heated:

NH 4 NO 2 \u003d 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 (31 ml of oxygen) dissolve in 1 liter of water. Therefore, in air dissolved in water, the content of oxygen 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 combines directly only with lithium:

6Li + N 2 = 2Li 3 N

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

3Ca + N 2 \u003d Ca 3 N 2, 2Al + N 2 \u003d 2AlN

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

N 2 + 3H 2 \u003d 2NH 3

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

Application. Nitrogen is used in large quantities to produce ammonia. It is widely used to create an inert environment - filling electric incandescent lamps and free space in mercury thermometers, when pumping flammable liquids. They nitride the surface of steel products, t. saturate their surface with nitrogen at high temperature. As a result, iron nitrides are formed in the surface layer, which give the steel greater hardness. 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, methods of obtaining. 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 the ammonium ion.

Ammonia - colorless gas with a characteristic odor, almost twice as light as air. When pressure is increased or cooled, it readily 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, the dissolved ammonia evaporates from the solution.

Chemical properties.

Interaction with acids:

NH 3 + HCl \u003d NH 4 Cl, NH 3 + H 3 PO 4 \u003d NH 4 H 2 PO 4

Interaction with oxygen:

4NH 3 + 3O 2 \u003d 2N 2 + 6H 2 O

Copper recovery:

3CuO + 2NH 3 \u003d 3Cu + N 2 + 3H 2 O

Receipt.

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

N 2 + 3H 2 \u003d 2NH 3

Application.

Liquid ammonia and its aqueous solutions are used as a 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. They are similar in structure 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 instance:

NH 3 + HNO 3 \u003d 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 \u003d NaCl + H 2 O + NH 3

2NH 4 Cl + H 2 SO 4 \u003d (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 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 products as a chemical baking powder, in dyeing fabrics, 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 a 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 unlimed, infertile soils. But these fertilizers also have their advantages: ammonium is much less subject to leaching, as it is fixed by soil particles and absorbed by microorganisms, and, in addition, the process of nitrofication occurs with it in the soil, i.e. conversion by microorganisms to 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 the 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 blue staining of 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 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, almost 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:

Interesting is the fact that with other 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, but requires incomparably higher temperatures.

Active metal nitrides are easily hydrolyzed by water:

As well as 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 ammonia yield 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 influence of an electric discharge or at a temperature of more than 2000 ° C and is reversible:

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

The 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 four-atomic P 4 molecules and is not a stable modification of phosphorus. Poisonous. At room temperature, it is soft and, like wax, can be easily cut with a knife. In air, it slowly oxidizes, 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 still more correct to consider its structure as molecular. Due to 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 it.

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

The interaction of phosphorus with non-metals

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

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

The reaction with halogens also proceeds intensively. So, during 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, it is possible to oxidize phosphorus with iodine only to an oxidation state of +3:

Unlike nitrogen phosphorus does not react with hydrogen.

The 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:

As well as aqueous solutions of non-oxidizing acids:

The 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 down the equations of such interactions for the Unified State Examination 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:

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 is [ 2 He] 2s 2 2p 3, characteristic oxidation states 0, -3, +3 and +5, less often +2 and +4 and the other state N v is considered to be relatively stable.

Nitrogen oxidation state 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 the 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 N≡N ˚σππ bond, which explains the chemical inertness of the element under normal conditions.

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

The main component of air is 78.09% by volume, 75.52 by mass. Nitrogen boils out of 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, N 2 reacts with fluorine and, to a very small extent, with oxygen:

N 2 + 3F 2 \u003d 2NF 3, N 2 + O 2 ↔ 2NO

The reversible reaction of obtaining ammonia proceeds 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 at Pt)

N 2 + 3H 2 ↔ 2NH 3 + 92 kJ

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

Nitrogen is chemically passive with respect to acids and alkalis, does not support combustion.

Receipt v industry- fractional distillation of liquid air or chemical removal of oxygen from air, for example, by the reaction 2C (coke) + O 2 \u003d 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 a switching reaction with moderate heating:

N -3 H 4 N 3 O 2 (T) \u003d 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

Binary compound, nitrogen oxidation state is - 3. A 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 nitrogen in the NH 3 molecule of a donor pair of electrons in the sp 3 hybrid orbital causes a characteristic addition reaction of a hydrogen cation, with the formation of a cation ammonium NH4. It liquefies under positive pressure at room temperature. In the liquid state, it is associated by hydrogen bonds. Thermally unstable. Let's well dissolve in water (more than 700 l/1 l of H 2 O at 20˚C); the proportion in the saturated solution is 34% by weight and 99% by volume, pH= 11.8.

Very reactive, prone to addition reactions. Burns in oxygen, reacts with acids. Shows 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. 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 \u003d Mg 3 N 2 +3 H 2 (600 ˚C)
NH 3 (g) + CO 2 (g) + H 2 O \u003d NH 4 HCO 3 (room temperature, pressure)
Receipt. V laboratories- displacement of ammonia from ammonium salts when heated with soda lime: Ca (OH) 2 + 2NH 4 Cl \u003d CaCl 2 + 2H 2 O + NH 3
Or boiling an aqueous solution of ammonia, followed by drying the gas.
In industry ammonia is produced 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 connection. White, in the crystal lattice - NH 3 and H 2 O molecules bound by a weak hydrogen bond. It is present in an aqueous solution of ammonia, a weak base (dissociation products are the NH 4 cation and the 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. It exhibits reducing properties (due to N -3) in a concentrated solution. It enters into the reaction of ion exchange and complex formation.

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 dissociation of the hydrate); thus, the ionic "ammonium hydroxide NH 4 OH" is practically not contained in the solution, there is no such compound in the solid hydrate either.
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 (diff.) = 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 a concentrated solution (18.5 - 25%) is an ammonia solution (produced by industry).

nitrogen oxides

nitrogen monoxideNO

Non-salt forming oxide. colorless gas. The radical contains a covalent σπ-bond (N꞊O), in the solid state an N 2 O 2 dimer with an N-N bond. Extremely thermally stable. Sensitive to atmospheric oxygen (turns brown). Slightly soluble in water and does not react with it. Chemically passive in relation to acids and alkalis. When heated, it reacts with metals and non-metals. highly reactive mixture of NO and NO 2 ("nitrous gases"). An intermediate product in the synthesis of nitric acid.
Equations of the most important reactions:
2NO + O 2 (ex.) = 2NO 2 (20˚C)
2NO + C (graphite) \u003d N 2 + CO 2 (400-500˚C)
10NO + 4P(red) = 5N 2 + 2P 2 O 5 (150-200˚C)
2NO + 4Cu \u003d N 2 + 2 Cu 2 O (500-600˚C)
Reactions to mixtures of NO and NO 2:
NO + NO 2 + H 2 O \u003d 2HNO 2 (p)
NO + NO 2 + 2KOH(razb.) \u003d 2KNO 2 + H 2 O
NO + NO 2 + Na 2 CO 3 \u003d 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 \u003d 3Hg 2 (NO 3) 2 + 2 NO+ 4 H 2 O
or reduction of nitrates:
2NaNO 2 + 2H 2 SO 4 + 2NaI \u003d 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, monomer NO 2 at room temperature, liquid colorless dimer N 2 O 4 (dianitrogen tetroxide) in the cold. Completely reacts with water, alkalis. Very strong oxidizing agent, corrosive to metals. It is used for the synthesis of nitric acid and anhydrous nitrates, as an oxidizer for rocket fuel, an oil cleaner from sulfur, and a catalyst for the oxidation of organic compounds. Poisonous.
The equation of the most important reactions:
2NO 2 ↔ 2NO + O 2
4NO 2 (l) + H 2 O \u003d 2HNO 3 + N 2 O 3 (syn.) (in the cold)
3 NO 2 + H 2 O \u003d 3HNO 3 + NO
2NO 2 + 2NaOH (diff.) \u003d NaNO 2 + NaNO 3 + H 2 O
4NO 2 + O 2 + 2 H 2 O \u003d 4 HNO 3
4NO 2 + O 2 + KOH \u003d 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 with atmospheric oxygen, in laboratories– interaction of concentrated nitric acid with reducing agents:
6HNO 3 (conc., mountains) + S \u003d H 2 SO 4 + 6NO 2 + 2H 2 O
5HNO 3 (conc.,hort.) + P (red) \u003d H 3 PO 4 + 5NO 2 + H 2 O
2HNO 3 (conc., mountains) + SO 2 \u003d H 2 SO 4 + 2 NO 2

dinitrogen oxideN 2 O

Colorless gas with a pleasant smell ("laughing gas"), N꞊N꞊О, formal nitrogen oxidation state +1, poorly soluble in water. Supports the 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 \u003d N 2 O + 2 H 2 O (195-245˚C)
used in medicine as an anesthetic.

dinitrogen trioxideN 2 O 3

At low temperatures, it is a blue liquid, ON꞊NO 2, the formal oxidation state of nitrogen is +3. At 20 ˚C, it decomposes by 90% into a mixture of colorless NO and brown NO 2 (“nitrous gases”, industrial smoke - “fox tail”). N 2 O 3 - acid oxide, forms HNO 2 with water in the cold, reacts differently when heated:
3N 2 O 3 + H 2 O \u003d 2HNO 3 + 4NO
With alkalis gives HNO 2 salts, for example NaNO 2 .
Obtained by the interaction of NO with O 2 (4NO + 3O 2 \u003d 2N 2 O 3) or with NO 2 (NO 2 + NO \u003d 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, 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 acidic oxide:
N 2 O 5 + H 2 O \u003d 2HNO 3
N 2 O 5 + 2NaOH \u003d 2NaNO 3 + H 2
Obtained by dehydration of fuming nitric acid:
2HNO 3 + P 2 O 5 \u003d N 2 O 5 + 2HPO 3
or oxidation of NO 2 with ozone at -78˚C:
2NO 2 + O 3 \u003d N 2 O 5 + O 2

Nitrites and nitrates

Potassium nitriteKNO 2 . White, hygroscopic. Melts without decomposition. Stable in dry air. Let's very well dissolve in water (forming colorless solution), it is hydrolyzed on anion. A typical oxidizing and reducing agent in an acidic environment, reacts very slowly in an alkaline environment. Enters into ion exchange reactions. Qualitative reactions on the NO 2 ion - discoloration of the purple solution of MnO 4 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 (t) + 2HNO 3 (conc.) \u003d NO 2 + NO + H 2 O + 2KNO 3
2KNO 2 (dil.) + O 2 (ex.) → 2KNO 3 (60-80 ˚C)
KNO 2 + H 2 O + Br 2 = KNO 3 + 2HBr
5NO 2 - + 6H + + 2MnO 4 - (violet) \u003d 5NO 3 - + 2Mn 2+ (bts.) + 3H 2 O
3 NO 2 - + 8H + + CrO 7 2- \u003d 3NO 3 - + 2Cr 3+ + 4H 2 O
NO 2 - (saturated) + NH 4 + (saturated) \u003d N 2 + 2H 2 O
2NO 2 - + 4H + + 2I - (BC) = 2NO + I 2 (black) ↓ = 2H 2 O
NO 2 - (razb.) + Ag + \u003d AgNO 2 (light yellow) ↓
Receipt 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 \u003d 2 KNO 2+ CaSO 4 (300 ˚C)

H itrat potassium KNO 3
technical name potassium, or indian salt , saltpeter. White, melts without decomposition, decomposes upon further heating. Air resistant. Highly soluble in water (high endo-effect, = -36 kJ), there is no hydrolysis. A strong oxidizing agent when fused (due to the release of atomic oxygen). In solution, it is reduced only by atomic hydrogen (in an acid 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 \u003d 2KNO 2 + O 2 (400-500 ˚C)

KNO 3 + 2H 0 (Zn, diluted 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 \u003d 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 (horizontal) + 4NO 2 + O 2 = 4KNO 3 + 2H 2 O

and in the lab:
KCl + AgNO 3 \u003d KNO 3 + AgCl ↓

Compounds with non-metals

All nitrogen halides NG 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 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 an unshared electron pair. To various chemical reagents and to heating, NF 3 is very stable.

The remaining 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 \u003d 4HCl + NCl 3

Nitrogen trichloride is a highly volatile (t bp = 71 degrees C) liquid with a pungent odor. A slight heating or impact is accompanied by an explosion with the release of a large amount of heat. In this case, NCl 3 decomposes into elements. The 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, they explode when heated and hit.

Connections with 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 \u003d 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 an atmosphere of nitrogen or ammonia. Oxides, halides and hydrides of transition metals can be used as starting materials:

2Ta + N 2 \u003d 2TaN; Mn 2 O 3 + 2NH 3 \u003d 2MnN + 3H 2 O

CrCl 3 + NH 3 = CrN + 3HCl; 2TiH 2 + 2NH 3 \u003d 2TiN + 5H 2

The use of nitrogen and nitrogen-containing compounds

The scope of nitrogen is very large - 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, 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 and hydrogen as ammonia is used in refrigeration plants, ammonia, circulating through a closed system of pipes, takes a large amount of heat during its evaporation. Potassium nitrate is used in the production of black powder, and gunpowder is used in hunting rifles, for exploration of ore minerals that occur underground. Smokeless gunpowder is obtained from pyroxylin, an ester of cellulose and nitric acid. Nitrogen-based organic explosives are used for tunneling in the mountains (TNT, nitroglycerin).