HNO 2 has a weak character. Very unstable, can only be found in dilute solutions:

2 HNO 2 NO + NO 2 + H 2 O.

Salts nitrous acid are called nitrites or nitrous acid. Nitrites are much more stable than HNO2, they are all toxic.

2HNO 2 + 2HI = I 2 + 2NO + 2H 2 O,

HNO 2 + H 2 O 2 = HNO 3 + H 2 O,

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

Structure of nitrous acid.

In the gas phase, the planar nitrous acid molecule exists in the form of two configurations: cis- and trans-:

At room temperature the trans isomer predominates: this structure is more stable. So, for cis - HNO2(G) DG° f= −42.59 kJ/mol, and for trans- HNO2(G) DG= −44.65 kJ/mol.

Chemical properties of nitrous acid.

In aqueous solutions there is an equilibrium:

When heated, the nitrous acid solution decomposes, releasing NO and the formation of nitric acid:

HNO2 V aqueous solutions dissociates ( K D=4.6 10 −4), slightly stronger acetic acid. Easily displaced by more strong acids from salts:

Nitrous acid exhibits oxidizing and reducing properties. Under the action of stronger oxidizing agents (hydrogen peroxide, chlorine, potassium permanganate), oxidation into nitric acid occurs:

In addition, it can oxidize substances that have reducing properties:

Preparation of nitrous acid.

Nitrous acid is obtained by dissolving nitric oxide (III) N2O3 in water:

In addition, it is formed when nitric oxide (IV) is dissolved in water. NO 2:

.

Application of nitrous acid.

Nitrous acid is used for diazotization of primary aromatic amines and the formation of diazonium salts. Nitrites are used in organic synthesis in the production of organic dyes.

Physiological effect of nitrous acid.

Nitrous acid is toxic and has a pronounced mutagenic effect, as it is a deaminating agent.

Nitrous acid exists either in solution or in the gas phase. It is unstable and, when heated, disintegrates into vapors:

2HNO 2 “NO+NO 2 +H 2 O

Aqueous solutions of this acid decompose when heated:

3HNO 2 “HNO 3 +H 2 O+2NO

This reaction is reversible, therefore, although the dissolution of NO 2 is accompanied by the formation of two acids: 2NO 2 + H 2 O = HNO 2 + HNO 3

Practically, by reacting NO 2 with water, HNO 3 is obtained:

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

By acidic properties Nitrous acid is only slightly stronger than acetic acid. Its salts are called nitrites and, unlike the acid itself, are stable. From solutions of its salts, a solution of HNO 2 can be obtained by adding sulfuric acid:

Ba(NO 2) 2 +H 2 SO 4 =2HNO 2 +BaSO 4 ¯

Based on data on its compounds, two types of structure of nitrous acid are suggested:

which correspond to nitrites and nitro compounds. Nitrites active metals have a type I structure, and low-active metals have a type II structure. Almost all salts of this acid are highly soluble, but silver nitrite is the most difficult. All salts of nitrous acid are poisonous. For chemical technology KNO 2 and NaNO 2 are important, which are necessary for the production of organic dyes. Both salts are obtained from nitrogen oxides:

NO+NO 2 +NaOH=2NaNO 2 +H 2 O or when heating their nitrates:

KNO 3 +Pb=KNO 2 +PbO

Pb is necessary to bind the released oxygen.

Of the chemical properties of HNO 2, oxidative ones are more pronounced, while it itself is reduced to NO:

However, many examples of such reactions can be given where nitrous acid exhibits reducing properties:

The presence of nitrous acid and its salts in a solution can be determined by adding a solution of potassium iodide and starch. Nitrite ion oxidizes iodine anion. This reaction requires the presence of H +, i.e. occurs in an acidic environment.

Nitric acid

In laboratory conditions, nitric acid can be obtained by the action of concentrated sulfuric acid on nitrates:

NaNO 3 +H 2 SO 4(k) =NaHSO 4 +HNO 3 The reaction occurs with low heating.

The production of nitric acid on an industrial scale is carried out by the catalytic oxidation of ammonia with atmospheric oxygen:

1. First, a mixture of ammonia and air is passed over a platinum catalyst at 800°C. Ammonia is oxidized to nitric oxide (II):

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

2. Upon cooling, further oxidation of NO occurs to NO 2: 2NO+O 2 =2NO 2

3. The resulting nitrogen oxide (IV) dissolves in water in the presence of excess O 2 to form HNO 3: 4NO 2 +2H 2 O+O 2 =4HNO 3

The starting products - ammonia and air - are thoroughly purified from harmful impurities, poisoning the catalyst (hydrogen sulfide, dust, oils, etc.).

The resulting acid is dilute (40-60% acid). Concentrated nitric acid (96-98% strength) is obtained by distilling dilute acid in a mixture with concentrated sulfuric acid. In this case, only nitric acid evaporates.

Physical properties

Nitric acid is a colorless liquid with a pungent odor. Very hygroscopic, “smoke” in air, because its vapors with air moisture form drops of fog. Mixes with water in any ratio. At -41.6°C it goes into a crystalline state. Boils at 82.6°C.

In HNO 3, the valency of nitrogen is 4, the oxidation state is +5. The structural formula of nitric acid is depicted as follows:

Both oxygen atoms associated only with nitrogen are equivalent: they are at the same distance from the nitrogen atom and each carry half the charge of an electron, i.e. the fourth part of nitrogen is divided equally between two oxygen atoms.

Electronic structure nitric acid can be derived as follows:

1. A hydrogen atom bonds with an oxygen atom by a covalent bond:

2. Due to the unpaired electron, the oxygen atom forms a covalent bond with the nitrogen atom:

3. Two unpaired electron nitrogen atoms form covalent bond with the second oxygen atom:

4. The third oxygen atom, when excited, forms a free 2p- orbital by electron pairing. The interaction of a nitrogen lone pair with a vacant orbital of the third oxygen atom leads to the formation of a nitric acid molecule:

Chemical properties

1. Dilute nitric acid exhibits all the properties of acids. It belongs to strong acids. Dissociates in aqueous solutions:

HNO 3 “Н + +NO - 3 Partially decomposes under the influence of heat and light:

4HNO 3 =4NO 2 +2H 2 O+O 2 Therefore, store it in a cool and dark place.

2. Nitric acid is characterized exclusively by oxidizing properties. The most important chemical property is interaction with almost all metals. Hydrogen is never released. The reduction of nitric acid depends on its concentration and the nature of the reducing agent. The degree of oxidation of nitrogen in the reduction products is in the range from +4 to -3:

HN +5 O 3 ®N +4 O 2 ®HN +3 O 2 ®N +2 O®N +1 2 O®N 0 2 ®N -3 H 4 NO 3

The reduction products from the interaction of nitric acid of different concentrations with metals of different activity are shown in the diagram below.

Concentrated nitric acid at ordinary temperatures does not interact with aluminum, chromium, and iron. It puts them into a passive state. A film of oxides forms on the surface, which is impermeable to concentrated acid.

3. Nitric acid does not react with Pt, Rh, Ir, Ta, Au. Platinum and gold dissolve in “royal vodka” - a mixture of 3 volumes of concentrated hydrochloric acid and 1 volume of concentrated nitric acid:

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

3Pt+4HNO 3 +12HCl=3PtCl 4 +4NO+8H 2 O 2HCl+PtCl 4 =H 2

The effect of “regia vodka” is that nitric acid oxidizes hydrochloric acid to free chlorine:

HNO 3 +HCl=Cl 2 +2H 2 O+NOCl 2NOCl=2NO+Cl 2 The released chlorine combines with metals.

4. Non-metals are oxidized with nitric acid to the corresponding acids, and depending on the concentration it is reduced to NO or NO 2:

S+bHNO 3(conc) =H 2 SO 4 +6NO 2 +2H 2 OP+5HNO 3(conc) =H 3 PO 4 +5NO 2 +H 2 O I 2 +10HNO 3(conc) =2HIO 3 +10NO 2 +4H 2 O 3P+5HNO 3(p asb) +2H 2 O= 3H 3 PO 4 +5NO

5. It also interacts with organic compounds.

Salts of nitric acid are called nitrates and are crystalline substances, highly soluble in water. They are obtained by the action of HNO 3 on metals, their oxides and hydroxides. Potassium, sodium, ammonium and calcium nitrates are called nitrates. Nitrate is used mainly as mineral nitrogen fertilizers. In addition, KNO 3 is used to prepare black powder (a mixture of 75% KNO 3, 15% C and 10% S). The explosive ammonal is made from NH 4 NO 3, aluminum powder and trinitrotoluene.

Salts of nitric acid decompose when heated, and the decomposition products depend on the position of the salt-forming metal in the series of standard electrode potentials:

Decomposition when heated (thermolysis) is an important property of nitric acid salts.

2KNO 3 =2KNO 2 +O 2

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

Salts of metals located in the series to the left of Mg form nitrites and oxygen, from Mg to Cu - metal oxide, NO 2 and oxygen, after Cu - free metal, NO 2 and oxygen.

Application

Nitric acid is the most important product of the chemical industry. Large quantities are spent on the preparation of nitrogen fertilizers, explosives, dyes, plastics, artificial fibers and other materials. Smoking

Nitric acid is used in rocket technology as a rocket fuel oxidizer.

Ammonium salts are very peculiar. All of them decompose easily, some spontaneously, for example ammonium carbonate:
(NH4)2CO3 = 2NH3 + H2O + CO2 (the reaction accelerates when heated).
Other salts, for example ammonium chloride (ammonia), sublimate when heated, i.e., they first decompose into ammonia and chloride under the influence of heating, and when the temperature decreases, ammonium chloride is formed again on the cold parts of the vessel:
heating
NH4Cl ⇄ NH3 + HCl
cooling
When heated, ammonium nitrate decomposes into nitrous oxide and water. This reaction can occur explosively:
NH4NO3 = N2O + H2O
Ammonium nitrite NH4NO2 decomposes when heated to form nitrogen and water, so it is used in the laboratory to obtain nitrogen.
When ammonium salts are exposed to alkalis, ammonia is released:
NH4Cl + NaOH = NaCl + NH3 + H2O
Ammonia release - characteristic feature for recognizing ammonium salts. All ammonium salts are complex compounds.

Ammonia and ammonium salts are widely used. Ammonia is used as a raw material for the production of nitric acid and its salts, as well as ammonium salts, which serve as good nitrogen fertilizers. Such fertilizers are ammonium sulfate (NH4)2SO4 and especially ammonium nitrate NH4NO3 or ammonium nitrate, the molecule of which contains two nitrogen atoms: one ammonium, the other nitrate. Plants first absorb ammonia and then nitrate. This conclusion belongs to the founder of Russian agrochemistry, Acad. D. N. Pryanishnikov, who devoted his works to plant physiology and substantiated the importance of mineral fertilizers in agriculture.
Ammonia in the form of ammonia is used in medicine. Liquid ammonia is used in refrigeration units. Ammonium chloride is used to make Leclanché dry galvanic cell. A mixture of ammonium nitrate with aluminum and coal, called ammonal, is a powerful explosive.
Ammonium carbonate is used in the confectionery industry as a leavening agent.

■ 25. On what property of ammonium carbonate is its use for loosening dough based?
26. How to detect ammonium ion in salt?
27. How to carry out a series of transformations:
N2 ⇄ NH3 → NO
↓
NH4N03

Oxygen compounds of nitrogen

It forms several compounds with oxygen, in which it exhibits different oxidation states.
There is nitrous oxide N2O, or, as it is called, “laughing gas”. It exhibits an oxidation state of + 1. In nitrogen oxide NO, nitrogen exhibits an oxidation state of + 2, in nitrous anhydride N2O3 - + 3, in nitrogen dioxide NO2 - +4, in nitrogen pentoxide, or nitric
anhydride, N2O5 - +5.
Nitrous oxide N2O is a non-salt-forming oxide. This is a gas that is quite soluble in water, but does not react with water. Nitrous oxide mixed with oxygen (80% N2O and 20% O2) produces a narcotic effect and is used for so-called gas anesthesia, the advantage of which is that it does not have a long aftereffect.
The rest of the nitrogen is highly poisonous. Their toxic effect usually occurs within a few hours after inhalation. First aid consists of ingesting a large amount of milk, inhaling pure oxygen, and resting the victim.

■ 28. List the possible oxidation states of nitrogen and corresponding to these oxidation states.
29. What first aid measures should be taken for poisoning with nitrogen oxides?

The most interesting and important nitrogen oxides are nitrogen oxide and nitrogen dioxide, which we will study.
Nitric oxide NO is formed from nitrogen and oxygen during strong electrical discharges. The formation of nitrogen oxide is sometimes observed in the air during a thunderstorm, but in very small quantities. Nitric oxide is a colorless, odorless gas. Nitric oxide is insoluble in water, so it can be collected above water in cases where the preparation is carried out in a laboratory. In the laboratory, nitric oxide is obtained from moderately concentrated nitric acid by its action on:
HNO3 + Cu → Cu(NO3)2 + NO + H2O
Arrange the coefficients in this equation yourself.
Nitric oxide can be produced in other ways, for example in a flame electric arc:
N2 + O2 ⇄ 2NO.
In the production of nitric acid, nitric oxide is obtained by the catalytic oxidation of ammonia, which was discussed in § 68, page 235.
Nitric oxide is a non-salt-forming oxide. It is easily oxidized by atmospheric oxygen and turns into nitrogen dioxide NO2. If oxidation is carried out in a glass vessel, colorless nitric oxide turns into a brown gas - nitrogen dioxide.

■ 30. When copper interacts with nitric acid, 5.6 liters of nitric oxide are released. Calculate how much copper reacted and how much salt was formed.

Nitrogen dioxide NO2 is a brown gas with a characteristic odor. It is highly soluble in water, as it reacts with water according to the equation:
3NO2 + H2O = 2HNO3 + NO
In the presence of oxygen, only nitric acid can be obtained:
4NO2 + 2H2O + O2 = 4HNO3
Molecules of nitrogen dioxide NO2 quite easily combine in pairs and form nitrogen tetroxide N2O4 - a colorless liquid, structural formula which

This process occurs in the cold. When heated, nitrogen tetroxide turns back into nitrogen dioxide.
Nitrogen dioxide is an acidic oxide because it can react with alkalis to form salt and water. However, due to the fact that nitrogen atoms in the N2O4 modification have a different number of valence bonds, when nitrogen dioxide reacts with alkali, two salts are formed - nitrate and nitrite:
2NO2 + 2NaOH = NaNO3 + NaNO2 + H2O
Nitrogen dioxide is obtained, as mentioned above, by oxidation of the oxide:
2NO + O2 = 2NO2
In addition, nitrogen dioxide is produced by the action of concentrated nitric acid on:
Сu + 4HNO3 = Cu(NO3)2 + 2NO2 + 2H2O
(conc.)
or better by calcining lead nitrate:
2Pb(NO3)2 = 2PbO + 4NO2 + O2

■ 31. List the methods for producing nitrogen dioxide, giving equations for the corresponding reactions.

32. Draw a diagram of the structure of the nitrogen atom in the +4 oxidation state and explain what its behavior should be in redox reactions.
33. 32 g of a mixture of copper and copper oxide were placed in concentrated nitric acid. The copper content in the mixture is 20%. What volume of what gas will be released? How many gram molecules of salt does this produce?

Nitrous acid and nitrites

Nitrous acid HNO2 is a very weak unstable acid. It exists only in dilute solutions (a = 6.3% in a 0.1 N solution). Nitrous acid easily decomposes to form nitrogen oxide and nitrogen dioxide
2HNO2 = NO + NO2 + H2O.
The oxidation state of nitrogen in nitrous acid is +3. With this degree of oxidation, we can conventionally assume that 3 electrons have been given up from the outer layer of the nitrogen atom and 2 valence electrons remain. In this regard, there are two possibilities for N+3 in redox reactions: it can exhibit both oxidizing and reducing properties, depending on which environment - oxidative or reducing - it enters.
Salts of nitrous acid are called nitrites. By treating nitrites with sulfuric acid, you can get nitrous acid:
2NaNO2 + H2SO4 = Na2SO4 + 2HNO2.
Nitrites are salts that are quite soluble in water. Like nitrous acid itself, nitrites can exhibit oxidizing properties when reacting with reducing agents, for example:
NaNO2 + KI + H2SO4 → I2 + NO…

Try to find the final products and arrange coefficients based on the electronic balance yourself.

Since the release is easy to detect using starch, this reaction can serve as a way to detect even small amounts of nitrite in drinking water, the presence of which is undesirable due to toxicity. On the other hand, nitrite nitrogen can be oxidized to N +5 under the action of a strong oxidizing agent.
NaNO2 + K2Cr2O7 + H2SO4 → NaNO3 + Cr2(SO4)3 + …

Find the rest of the reaction products yourself, compose electronic balance and set the coefficients.

■ 34. Complete the equation.
HNO2 + KMnO4 + H2SO4 → … (N +5, Mn +2).
35. List the properties of nitrous acid and nitrites.

Nitric acid

HNO3 is a strong electrolyte. This is a volatile liquid. Pure boils at a temperature of 86°, has no color; its density is 1.53. Laboratories typically receive 65% HNO3 with a density of 1.40.
smokes in the air, since its vapors, rising into the air and combining with water vapor, form droplets of fog. Nitric acid mixes with water in any ratio. It has a pungent odor and evaporates easily, so concentrated nitric acid should only be poured under pressure. If it comes in contact with skin, nitric acid can cause severe burns. A small burn makes itself known as a characteristic yellow spot on the skin. Severe burns can cause ulcers. If nitric acid comes into contact with the skin, it should be quickly washed off with plenty of water and then neutralized with a weak solution of soda.

Concentrated 96-98% nitric acid rarely enters the laboratory and during storage quite easily, especially in light, it decomposes according to the equation:
4HNO3 = 2H2O + 4NO2 + O2
It is permanently colored with nitrogen dioxide yellow. Excess nitrogen dioxide gradually evaporates from the solution, accumulates in the solution, and the acid continues to decompose. In this regard, the concentration of nitric acid gradually decreases. At a concentration of 65%, nitric acid can be stored for a long time.
Nitric acid is one of the strongest oxidizing agents. It reacts with almost all metals, but without releasing hydrogen. The pronounced oxidizing properties of nitric acid have a so-called passivating effect on some (,) compounds. This is especially true for concentrated acids. When exposed to it, a very dense acid-insoluble oxide film is formed on the metal surface, protecting the metal from further exposure to acid. The metal becomes "passive". .
However, nitric acid reacts with most metals. In all reactions with metals, nitrogen is reduced in nitric acid, and the more completely, the more dilute the acid and the more active the metal.

The concentrated acid is reduced to nitrogen dioxide. An example of this is the reaction with copper given above (see § 70). Dilute nitric acid with copper is reduced to nitric oxide (see § 70). More active ones, for example, reduce dilute nitric acid to nitrous oxide.
Sn + HNO3 → Sn(NO3)2 + N2O
With very strong dilution with an active metal, for example zinc, the reaction reaches the formation of an ammonium salt:
Zn + HNO3 → Zn(NO3)2 + NH4NO3

In all the given reaction schemes, arrange the coefficients by creating an electronic balance yourself.

■ 36. Why does the concentration of nitric acid decrease when stored in the laboratory, even in well-sealed containers?
37. Why does concentrated nitric acid have a yellowish-brown color?
38. Write the equation for the reaction of dilute nitric acid with iron. The reaction products are iron(III) nitrate, and a brown gas is released.
39. Write down in your notebook all the reaction equations that characterize the interaction of nitric acid with metals. List what metal nitrates, in addition to metal nitrates, are formed in these reactions.

Many can burn in nitric acid, such as coal and:
C + HNO3 → NO + CO2
P + HNO3 → NO + H3PO4

The free one is oxidized to phosphoric acid. when boiled in nitric acid, it turns into S+6 and from free sulfur is formed:
HNO3 + S → NO + H2SO4

Complete the reaction equations yourself.

Complex ones can also burn in nitric acid. For example, turpentine and heated sawdust burn in nitric acid.
Nitric acid can also oxidize hydrochloric acid. A mixture of three parts hydrochloric acid and one part nitric acid is called aqua regia. This name is given because this mixture also oxidizes platinum, which is not affected by any acids. The reaction proceeds in the following stages: in the mixture itself, the chlorine ion is oxidized into a free one and nitrogen is reduced to form nitrosyl chloride:
HNO3 + 3HCl ⇄ Cl2 + 2H2O + NOCl
aqua regia nitrosyl chloride
The latter easily decomposes into nitric oxide and is free according to the equation:
2NOCl = 2NO + Cl2
Metal placed in aqua regia is easily oxidized by nitrosyl chloride:
Au + 3NOCl = AuCl3 + 3NO
Nitric acid can react with nitration with organic substances. In this case, a concentrated one must be present. A mixture of concentrated nitric and sulfuric acids is called a nitrating mixture. Using such a mixture, nitroglycerin can be obtained from glycerin, nitrobenzene from benzene, nitrocellulose from fiber, etc. In a highly diluted state, nitric acid exhibits the characteristic properties of acids.

■ 40. Give your own examples of typical properties of acids in relation to nitric acid. Write the equations in molecular and. ionic forms.
41. Why are bottles of concentrated nitric acid prohibited from being transported packed in wood shavings?
42. When concentrated nitric acid is tested with phenolphthalein, phenolphthalein acquires an orange color rather than remaining colorless. What explains this?

It is very easy to obtain nitric acid in the laboratory. It is usually obtained by displacing its salts with sulfuric acid, for example:
2KNO3 + H2SO4 = K2SO4 + 2HNO3
In Fig. 61 shows a laboratory installation for the production of nitric acid.
In industry, ammonia is used as a raw material for the production of nitric acid. As a result of the oxidation of ammonia in the presence of a platinum catalyst, nitrogen oxide is formed:
4NH3 + 5O2 = 4NO + 6H2O
As stated above, nitric oxide is easily oxidized by atmospheric oxygen into nitrogen dioxide:
2NO + O2 = 2NO2
and nitrogen dioxide, combining with water, forms nitric acid and again nitric oxide according to the equation:

3NO2 + H2O = 2HNO3 + NO.
Nitric oxide is then supplied again for oxidation:
The first stage of the process - the oxidation of ammonia into nitrogen oxide - is carried out in a contact apparatus at a temperature of 820°. The catalyst is a grid of platinum with an admixture of rhodium, which is heated before starting the apparatus. Since the reaction is exothermic, the grids are subsequently heated due to the heat of the reaction itself. The nitrogen oxide released from the contact apparatus is cooled to a temperature of about 40°, since the oxidation process of nitrogen oxide proceeds faster at a lower temperature. At a temperature of 140°, the resulting nitrogen dioxide decomposes again into oxides of nitrogen and oxygen.

The oxidation of nitrogen oxide into dioxide is carried out in towers called absorbers, usually under a pressure of 8-10 atm. They simultaneously absorb (absorb) the resulting nitrogen dioxide with water. To better absorb nitrogen dioxide, the solution is cooled. The result is 50-60% nitric acid.
The concentration of nitric acid is carried out in the presence of concentrated sulfuric acid in distillation columns. forms hydrates with the available water with a boiling point higher than that of nitric acid, so nitric acid vapors are quite easily released from the mixture. By condensing these vapors, 98-99% nitric acid can be obtained. Typically, a more concentrated acid is rarely used.

■ 43. Write down in your notebook all the equations of the reactions that occur when producing nitric acid by laboratory and industrial methods.
44. How to carry out a series of transformations:

45. How much of a 10% solution can be prepared from nitric acid obtained by reacting 2.02 kg of potassium nitrate with an excess of sulfuric acid?
46. ​​Determine the molarity of 63% nitric acid.
47. How much nitric acid can be obtained from 1 ton of ammonia at 70% yield?
48. The cylinder was filled with nitric oxide by displacing water. Then, without removing it from the water, a tube from a gasometer was placed under it.
(see Fig. 34) and began to skip. Describe what should be observed in the cylinder if excess oxygen was not allowed. Justify your answer with reaction equations.

Rice. 62. Combustion of coal in molten saltpeter. 1 - molten saltpeter; 2 - burning coal; 3 - sand.

Nitric acid salts

Salts of nitric acid are called nitrates. Nitrates alkali metals, as well as calcium and ammonium are called nitrate. For example, KNO3 is potassium nitrate, NH4NO3 is ammonium nitrate. Natural deposits of sodium nitrate are found in huge quantities in Chile, which is why this salt is called Chilean nitrate.

Rice. 62. Burning coal in molten saltpeter. 1 - molten saltpeter; 2 - burning coal; 3 - sand.

Salts of nitric acid, like itself, are strong oxidizing agents. For example, alkali metal salts are separated during melting according to the equation:

2KNO3 = 2KNO2+ O2

Thanks to this, coal and other flammable substances burn in molten saltpeter (Fig. 62).
Salts heavy metals also decompose with the release of oxygen, but according to a different pattern.
2Pb(NO3)2 = 2PbO + 4NO2 + O2

Rice. 63. Nitrogen cycle in nature

Potassium nitrate is used to make black gunpowder. To do this, it is mixed with coal and sulfur. It is not used for this purpose as it is hygroscopic. When ignited, black powder burns intensely according to the equation:
2KNO3 + 3С + S = N2 + 3CO2 + K2S
Calcium and ammonium nitrates are very good nitrogen fertilizers. Recently, potassium nitrate has also become widespread as a fertilizer.
Nitric acid is widely used in the production of chemical pharmaceuticals (streptocide), organic dyes, celluloid, film and photographic films. Salts of nitric acid are widely used in pyrotechnics.
In nature, there is a nitrogen cycle in which plants, when they die, return the nitrogen obtained from it back to the soil. Animals, feeding on plants, return nitrogen to the soil in the form of feces, and after death, their corpses rot and thereby also return the nitrogen received from it to the soil (Fig. 63). By harvesting a crop, a person interferes with this cycle, disrupts it and thereby depletes the soil of nitrogen, so it is necessary to apply nitrogen to the fields in the form of mineral fertilizers.

■ 49. How to carry out a series of transformations

Nitrous acid

If you heat potassium or sodium nitrate, they lose some of the oxygen and turn into salts of nitrous acid HNO2. Decomposition is easier in the presence of lead, which binds the released oxygen:

Nitrous acid salts - nitrites - form crystals that are highly soluble in water (with the exception of silver nitrite). Sodium nitrite NaNO 2 is used in the production of various dyes.

When a solution of some nitrite is exposed to dilute sulfuric acid, free nitrous acid is obtained:

It is one of the weak acids (K=A- 10~ 4) and is known only in highly dilute aqueous solutions. When the solution is concentrated or heated, nitrous acid decomposes:

The oxidation degree of nitrogen in nitrous acid is +3, i.e. is intermediate between the lowest and highest possible values ​​of the degree of nitrogen oxidation. Therefore, HNO 2 exhibits redox duality. Under the influence of reducing agents it is reduced (usually to NO), and in reactions with oxidizing agents it is oxidized to HNO 3. Examples include the following reactions:

Nitric acid

Pure nitric acid HNO3 is a colorless liquid with a density of 1.51 g/cm3, which solidifies into a transparent crystalline mass at -42 0C. In the air, it “smoke”, like concentrated hydrochloric acid, since its vapors form small droplets of fog with the moisture in the air.

Nitric acid is not strong. Already under the influence of light it gradually decomposes:

The higher the temperature and the more concentrated the acid, the faster the decomposition occurs. The released nitrogen dioxide dissolves in the acid and gives it a brown color.

Nitric acid is one of the most powerful acids; in dilute solutions it completely decomposes into H + and NO 3 ions.

Characteristic property nitric acid is its pronounced oxidizing ability. Nitric acid is one of the most energetic oxidizing agents. Many non-metals are easily oxidized by it, turning into the corresponding acids. Thus, sulfur, when boiled with nitric acid, gradually oxidizes into sulfuric acid, phosphorus - into phosphorus. A smoldering coal immersed in concentrated HNO 3 flares up brightly.

Nitric acid acts on almost all metals (with the exception of gold, platinum, tantalum, rhodium, iridium), turning them into nitrates, and some metals into oxides.

Concentrated HNO 3 passivates some metals. Lomonosov also discovered that iron, which easily dissolves in dilute nitric acid, does not dissolve in cold concentrated HNO 3. Later it was found that nitric acid has a similar effect on chromium and aluminum. These metals pass under the influence of concentrated nitric acid into a passive state (see § 100).

The oxidation degree of nitrogen in nitric acid is +5. Acting as an oxidizing agent, HNO 3 can be reduced to various products:

Which of these substances is formed, i.e. how deeply nitric acid is reduced in any given case depends on the nature of the reducing agent and the reaction conditions, primarily on the concentration of the acid. The higher the concentration of HNO 3, the less deeply it is reduced. In reactions with concentrated acid NO 2 is most often released. When dilute nitric acid reacts with low-active metals, such as copper, NO is released. In the case of more active metals - iron, zinc - N 2 O is formed. Highly diluted nitric acid reacts with active metals - zinc, magnesium, aluminum - to form ammonium ion, which gives ammonium nitrate with the acid. Usually several products are formed simultaneously.

For illustration, here are the reaction schemes for the oxidation of some metals with nitric acid:

When nitric acid acts on metals, hydrogen, as a rule, is not released.

When non-metals are oxidized, concentrated nitric acid, as in the case of metals, is reduced to NO 2, for example:

A more dilute acid is usually reduced to NO, for example:

The given diagrams illustrate the most typical cases of interaction of nitric acid with metals and non-metals. In general, redox reactions involving HNO 3 are complex.

A mixture consisting of 1 volume of nitric and 3-4 volumes of concentrated hydrochloric acid is called royal vodka. Aqua regia dissolves some metals that do not react with nitric acid, including the “king of metals” - gold. Its action is explained by the fact that nitric acid oxidizes hydrochloric acid, releasing free chlorine and forming nitrogen chloroxide(III), or nitrosyl chloride, NOCl:

Nitrosyl chloride is intermediate product reactions and decomposes:

Chlorine at the moment of release consists of atoms, which determines the high oxidizing ability of aqua regia. The oxidation reactions of gold and platinum proceed mainly according to the following equations:

With an excess of hydrochloric acid, gold (III) chloride and platinum (IV) chloride form complex compounds H[AuC1 4 ] and H 2 .

For many organic matter nitric acid acts so that one or more hydrogen atoms in the molecule organic compound are replaced by nitro groups - NO 2. This process is called nitration and has great value in organic chemistry.

The electronic structure of the HNO 3 molecule is discussed in § 44.

Nitric acid is one of the the most important connections nitrogen: it is used in large quantities in the production of nitrogen fertilizers, explosives and organic dyes, serves as an oxidizing agent in many chemical processes, is used in the production of sulfuric acid using the nitrous method, and is used for the manufacture of cellulose varnishes and film.

Salts of nitric acid are called nitrates. All of them dissolve well in water, and when heated, they decompose, releasing oxygen. In this case, the nitrates of the most active metals turn into nitrites:

Nitrates of most other metals decompose when heated into metal oxide, oxygen and nitrogen dioxide. For example:

Finally, nitrates of the least active metals (for example, silver, gold) decompose when heated to the free metal:

Easily splitting off oxygen, nitrates are energetic oxidizing agents at high temperatures. Their aqueous solutions, on the contrary, exhibit almost no oxidizing properties.

The most important are sodium, potassium, ammonium and calcium nitrates, which in practice are called saltpeter.

Sodium nitrate NaNO3, or sodium nitrate, sometimes also called Chilean saltpeter, is found in large quantities found in nature only in Chile.

Potassium nitrate KNO3, or potassium nitrate, is also found in nature in small quantities, but is mainly obtained artificially by reacting sodium nitrate with potassium chloride.

Both of these salts are used as fertilizers, and potassium nitrate contains two elements necessary for plants: nitrogen and potassium. Sodium and potassium nitrates are also used in glass melting and food industry for canning food.

Calcium nitrate Ca(NO 3) 2, or calcium nitrate, obtained in large quantities by neutralizing nitric acid with lime; used as fertilizer.

Ammonium nitrate NH4NO3.

• The student is encouraged to create his own complete equations these reactions.

HNO2 Physical properties State hard Molar mass 47.0134 g/mol Density 1.685 (liquid) Thermal properties T. float. 42.35 °C T. kip. 158 °C Chemical properties pK a 3.4 Solubility in water 548 g/100 ml Classification Reg. CAS number Data given is based on standard conditions (25 °C, 100 kPa) unless otherwise stated.

Nitrous acid HNO 2 is a weak monobasic acid that exists only in dilute aqueous solutions, colored faint blue, and in the gas phase. Salts of nitrous acid are called nitrites or nitrous acids. Nitrites are much more stable than HNO 2, all of which are toxic.

Structure

In the gas phase, the planar nitrous acid molecule exists in two configurations cis- And trance-.

cis isomer	trans isomer

At room temperature, the trans isomer predominates: this structure is more stable. Thus, for cis-HNO 2 (g) DG° f = −42.59 kJ/mol, and for trans-HNO 2 (g) DG = −44.65 kJ/mol.

Chemical properties

In aqueous solutions there is an equilibrium:

$\backslash mathsf(2HNO\_2\; \backslash rightleftarrows\; N\_2O\_3\; +\; H\_2O\; \backslash rightleftarrows\; NO\; \backslash uparrow\; +\; NO\_2\; \backslash uparrow\; +\; H\_2O)$

When the solution is heated, nitrous acid decomposes to release and form nitric acid:

$\backslash mathsf(3HNO\_2\; \backslash rightleftarrows\; HNO\_3\; +\; 2NO\; \backslash uparrow\; +\; H\_2O)$

HNO 2 is a weak acid. It dissociates in aqueous solutions (K D =4.6·10−4), slightly stronger than acetic acid. Easily replaced by stronger acids from salts:

$\backslash mathsf(H\_2SO\_4\; +\; 2NaNO\_2\; \backslash rightarrow\; Na\_2SO\_4\; +\; 2HNO\_2)$

Nitrous acid exhibits both oxidizing and reducing properties. Under the action of stronger oxidizing agents (hydrogen peroxide, chlorine, potassium permanganate) it is oxidized into nitric acid:

$\backslash mathsf(HNO\_2\; +\; H\_2O\_2\; \backslash rightarrow\; HNO\_3\; +\; H\_2O)$ $\backslash mathsf(HNO\_2\; +\; Cl\_2\; +\; H\_2O\backslash rightarrow\; HNO\_3\; +\; 2HCl)$ $\backslash mathsf(5HNO\_2\; +\; 2KMnO\_4\; +\; HNO\_3\; \backslash rightarrow\; 2Mn(NO\_3)\_2\; +\; 2KNO\_3\; +\; 3H\_2O)$

At the same time, it is capable of oxidizing substances with reducing properties:

\mathsf(2HNO_2 + 2HI \rightarrow 2NO\uparrow + I_2 +2H_2O)

Receipt

Nitrous acid can be obtained by dissolving nitric oxide (III) N 2 O 3 in water:

$\backslash mathsf(N\_2O\_3\; +\; H\_2O\; \backslash rightarrow\; 2HNO\_2)$ $\backslash mathsf(2NO\_2\; +\; H\_2O\; \backslash rightarrow\; HNO\_3\; +\; HNO\_2)$

Application

Nitrous acid is used to diazotize primary aromatic amines and form diazonium salts. Nitrites are used in organic synthesis in the production of organic dyes.

Physiological action

Nitrous acid is toxic and has a pronounced mutagenic effect, since it is a deaminating agent.

Sources

• Karapetyants M. Kh., Drakin S. I. General and inorganic chemistry. M.: Khimiya1994

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Links

• // Encyclopedic Dictionary of Brockhaus and Efron: in 86 volumes (82 volumes and 4 additional). - St. Petersburg. , 1890-1907.

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Excerpt characterizing Nitrous Acid

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