Which reaction is qualitative for ammonium salts. Ammonia. Ammonium salts. Let us recall the general properties of salts

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
The release of ammonia is a characteristic sign for the recognition of 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 an electric arc flame:
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 the nitrogen atoms in the N2O4 modification have different number 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 nitrites 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 influence of a strong oxidizing agent.
NaNO2 + K2Cr2O7 + H2SO4 → NaNO3 + Cr2(SO4)3 + …

Find the remaining reaction products yourself, draw up an electronic balance and arrange 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 which metals, 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 organic substances. In this case, concentrated 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 characteristic properties 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.
Then the nitric oxide is again supplied 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. Usually more concentrated acid 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 with a 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 of 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. IN Lately became widespread as a fertilizer and potassium nitrate.
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 they receive 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

Technical nitric acid The production of nitric acid is carried out in three ways, which we will describe in the order in which they began to be used...

CHECKING COMPLETION OF TASKS AND ANSWERS TO QUESTIONS 4. In order to answer these questions, carefully read §...

Ammonium salts When neutralized ammonia solutions are evaporated, ammonium ions combine with the anions of the taken acids, forming solid crystalline substances, possessing ionic...





















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Lesson type: lesson using multimedia

Lesson objectives:

  • Educational: Systematize students’ knowledge about salts; formation at the interdisciplinary level of a system of knowledge about ammonium salts, which are of great practical importance.
  • Educational: develop practical skills in conducting qualitative reactions to ammonium salts; the ability to analyze what you see; development logical thinking
  • ; development of cognitive interest when performing theoretical and practical tasks. Educational:

Increase cognitive activity and activity of students; developing the ability to work in a team.

Lesson equipment and reagents:

1. On the teacher’s desk: solution of hydrochloric acid HCl; ammonium hydroxide NH 4 OH; ammonium chloride NH 4 Cl; sodium chloride NaCl; water H 2 O; litmus; phenolphthalein; Ammonium bichromate (NH 4) 2 Cr 2 O 7. 2. On the students’ table: ammonium sulfate (NH 4) 2 SO 4; sulfuric acid

3.H2SO4; barium chloride BaCl; ammonium chloride NH 4 Cl; sodium hydroxide NaOH;

Interactive board.

Qualitative reactions to ammonium salts. Basic concepts:

Ammonium cation, ammonium salts.

During the classes

1. Organizational moment

The teacher checks readiness for the lesson and announces the topic of the lesson. 2. Check background knowledge

(10 min)

There are substances (salts) in glass containers on the teacher’s desk.

Teacher: This container contains an amazing substance. It was once considered “the grace of God,” a symbol of well-being.

:. But it can also destroy living things; because of it, even the sea can become dead.

However, it is difficult to list where it is used.

What is in this bottle?

(Suggested answer is salt.)

What substances do we classify as salts?

Training exercise:

Select salts from the listed substances and name them: KCl NaOH KOH
CO2 H2SO4 Ba(NO3)2 CuSO4
MgO NH4Cl H2S AgNO3

(NH 4) 2 SO 4

Teacher: What unusual salts have you come across?

These salts contain a complex cation - ammonium cation.

The teacher asks the class to formulate the concept of ammonium salt (salts consisting of ammonium cations and anions of an acid residue).

Teacher: Where do you think these salts can be used? Why?

Students: In agriculture, because they contain a vital element for plants - nitrogen.

For a comprehensive description of the practical significance of ammonium salts, it is advisable to listen to a mini report from students.

3. Learning new material (15 min)

What's common in physical properties salts and ammonium salts?

Students: solid, white, crystalline substances, highly soluble in water, electrolytes.

Teacher: Check the solubility of ammonium sulfate in practice, write down the physical properties in your notebook ( conducting an experiment).

Teacher: Let's consider the chemical properties of ammonium salts.

Let's remember general properties salts:

(The notebook contains reference summary).

A) dissociation - let’s write the equations for the dissociation of salts

  • Ammonium chloride
  • Ammonium sulfate

B) Interaction with acids

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

Carry out the experiment, indicate the signs of the reaction ( conducting an experiment).

B) Interaction with alkalis.

Place phenolphthalein paper into a test tube and observe a color change ( conducting an experiment)

Conclusion: this reaction is qualitative for ammonium salts

D) Interaction with salts

NH 4 Cl + AgNO 3 NH 4 NO 3 + AgCl

Students perform an experiment, write down the reaction equation and check them independently with the equation on the slide ( conducting an experiment).

Creative task: Determine which of the three test tubes contains ammonium sulfate. Write down the equations of the reactions performed ( conducting an experiment).

Specific properties of salts:

D) Decomposition of ammonium salts

Demonstration experiment: decomposition of ammonium dichromate; ammonium chloride:

NH 4 Cl NH 3 + HCl

(NH 4) 2 Cr 2 O 7 N 2 + Cr 2 O 3 +4H 2 O

E) Hydrolysis of ammonium salts

NH 4 Cl + H 2 O NH 4 OH + HCl

Demonstration experience.

Conclusion: the medium is alkaline, litmus is blue, phenolphthalein is crimson.

Teacher: Remember in what ways salts can be obtained.

Students: When a base and an acid interact; salts and salts; salts and acids.

The teacher demonstrates experiments, students write down a supporting summary using a slide.

A) NH 4 OH + HCl NH 4 Cl + H 2 O

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

B) (NH 4) 2 CO 3 + 2HCl 2 NH 4 Cl + H 2 O + CO 2

Conclusion: when writing equations, it is necessary to comply with the condition of irreversibility of chemical reactions.

4. Consolidation of the material covered (15 min)

Exercise No. 1.

Select and name ammonium salts:

Option I Option II
NaNO3 BaCl2

Peer testing in pairs.

Self-test.

NH 3 + HCl NH 4 Cl

2NH 3 + H 2 SO 4 (NH 4) 2 SO 4c) 3

React with sulfuric acid

The correct answer on the slide is marked with dots in the picture, then the dots are connected to form a smile.

Students compare their answers with the example on the screen and try to independently analyze the mistakes made. The teacher corrects the students' answers.

Exercise No. 4.(A game exercise that develops the children’s desire to find the correct answer, after which they can open the safe).

"Golden Key"

Determine the safe code.

List the sequence of numbers (in ascending order) that determine the properties of ammonium sulfate.

Test:

  1. Dissolves in water.
  2. Not electrolytes.
  3. White crystalline substance.
  4. Smells like ammonia.
  5. Reacts with barium chloride.
  6. Reacts with calcium hydroxide.
  7. Decomposes when heated.
  8. I DON'T KNOW THE WORDS
  9. reacts with orthophosphoric acid.
  10. Reacts with I DON'T KNOW THE NAME

Answer: 1345678

The teacher asks the children to write down possible reaction equations.

5. Homework

Creative task: Salt white, soluble in water, with silver nitrate it forms a white cheesy precipitate, the combustion of which produces nitrogen. Name the salt, write the reaction equations in molecular and ionic form.

6. Lesson conclusions (3min)

Completing the lesson in an interesting, creative form will put every child in a good mood and increase the quality of learning the material learned in the lesson.

Composing a cinquain (an interesting, non-rhyming poem that requires information in clear terms, which allows you to describe what you saw and heard):

Students write a syncwine, working in pairs, using the textbook and supporting notes.

  • Ammonium salts
  • Useful, important
  • Transform deserts into oases
  • React like all salts with salts, acids and alkalis
  • They decompose and are used in agriculture.

7. Summing up the lesson, the teacher quotes the words:“The thinking mind does not feel happy until it manages to connect disparate facts into one” (Hevelsey).

Nitrogen forms several compounds with hydrogen; of them highest value has ammonia - a colorless gas with a characteristic pungent odor (the smell of “ammonia”).

In the laboratory, ammonia is usually produced by heating ammonium chloride with slaked lime. The reaction is expressed by the equation

The released ammonia contains water vapor. To dry it, it is passed through soda lime (a mixture of lime and caustic soda).

Rice. 114. A device for demonstrating the combustion of ammonia in oxygen.

The mass of 1 liter of ammonia under normal conditions is 0.77 g. Since this gas is much lighter than air, it can be collected in vessels turned upside down.

When cooled to ammonia under normal pressure it turns into a clear liquid that solidifies at .

The electronic structure and spatial structure of the ammonia molecule are discussed in § 43. In liquid ammonia, the molecules are connected to each other by hydrogen bonds, which determines the relatively high boiling point of ammonia, which does not correspond to its low molecular weight (17).

Ammonia is very soluble in water: 1 volume of water dissolves at room temperature about 700 volumes of ammonia. The concentrated solution contains (mass) and has a density of . A solution of ammonia in water is sometimes called ammonia. Regular medical ammonia contains. As the temperature increases, the solubility of ammonia decreases, so it is released from a concentrated solution when heated, which is sometimes used in laboratories to obtain small quantities of ammonia gas.

At low temperatures, a crystalline hydrate can be isolated from an ammonia solution, melting at -. A crystalline hydrate of the composition is also known. In these hydrates, water and ammonia molecules are connected to each other by hydrogen bonds.

Chemically, ammonia is quite active; it interacts with many substances. In ammonia, nitrogen has the lowest oxidation state. Therefore, ammonia has only reducing properties. If a current is passed through a tube inserted into another wide tube (Fig. 114), through which oxygen passes, the ammonia can be easily ignited; it burns with a pale greenish flame. When ammonia burns, water and free nitrogen are formed:

Under other conditions, ammonia can be oxidized to nitrogen oxide (see § 143).

Unlike hydrogen compounds of non-metals of groups VI and VII, ammonia does not have acidic properties. However, hydrogen atoms in its molecule can be replaced by metal atoms.

When hydrogen is completely replaced by a metal, compounds called nitrides are formed. Some of them, such as calcium and magnesium nitrides, are obtained by the direct reaction of nitrogen with metals at high temperatures;

When in contact with water, many nitrides completely hydrolyze to form ammonia and metal hydroxide. For example:

When only one hydrogen atom in ammonia molecules is replaced by metals, metal amides are formed. Thus, by passing ammonia over molten sodium, sodium amide can be obtained in the form of colorless crystals:

Water decomposes sodium amide;

Possessing strong basic and water-removing properties, sodium amide has found use in some organic syntheses, for example, in the production of indigo dye and some drugs.

Hydrogen in ammonia can also be replaced by halogens. Thus, the action of chlorine on a concentrated solution of ammonium chloride produces chlorine nitride, or nitrogen chloride,

in the form of a heavy oily explosive liquid.

Iodine nitride (nitrogen iodide), which is formed in the form of a black, water-insoluble powder when iodine reacts with ammonia, has similar properties. When wet it is safe, but when dried it explodes at the slightest touch; in this case, violet iodine vapor is released.

With fluorine, nitrogen forms stable nitrogen fluoride.

From the data in table. 6 (p. 118) it is clear that the electronegativity of chlorine and sodium is less, and fluorine is greater, than the electronegativity of nitrogen. It follows that in compounds and the oxidation degree of nitrogen is -3, and in it is equal to . Therefore, nitrogen fluoride differs in properties from chlorine and iodine nitrides. For example, when interacting with water, ammonia is formed, and in this case, nitrogen oxide (III) is obtained;

The nitrogen atom in the ammonia molecule is connected by three covalent bonds to hydrogen atoms and retains one lone pair of electrons:

Acting as a donor of an electron pair, the nitrogen atom can participate in the formation of the fourth covalent bond with other atoms or ions that have electron-withdrawing properties.

This explains the extremely characteristic ability of ammonia to enter into addition reactions.

Examples of complex compounds formed by ammonia as a result of addition reactions are given in and 201, as well as in Chap. XVIII. Above (p. 124) the interaction of a molecule with a hydrogen ion, leading to the formation of ammonium ion, has already been considered:

In this reaction, ammonia serves as a proton acceptor and, therefore, from the point of view of the proton theory of acids and bases (p. 237), exhibits the properties of a base. Indeed, when reacting with acids that are in a free state or in solution, ammonia neutralizes them, forming ammonium salts. For example, with hydrochloric acid ammonium chloride is obtained:

The interaction of ammonia with water also leads to the formation of not only ammonia hydrates, but also partially ammonium ions:

As a result, the concentration of ions in the solution increases. This is why aqueous solutions of ammonia have an alkaline reaction. However, according to established tradition, an aqueous solution of ammonia is usually designated by the formula and called ammonium hydroxide, and the alkaline reaction of this solution is considered as the result of the dissociation of molecules.

Ammonia is a weak base. At the equilibrium constant of its ionization (see the previous equation) is equal to . A one-molar aqueous solution of ammonia contains only 0.0042 equivalents of and ions; such a solution has .

Most ammonium salts are colorless and highly soluble in water. In some of their properties they are similar to salts of alkali metals, especially potassium (the ions have similar sizes).

Since an aqueous solution of ammonia is a weak base, ammonium salts in solutions hydrolyze. Solutions of salts formed by ammonia and strong acids have a slightly acidic reaction.

Ammonium ion hydrolysis is usually written in this form:

However, it is more correct to consider it as a reversible transition of a proton from an ammonium ion to a water molecule:

When an alkali is added to an aqueous solution of any ammonium salt, the ions are bound by OH- ions into water molecules and the hydrolysis equilibrium shifts to the right. The process that occurs can be expressed by the equation:

When the solution is heated, ammonia evaporates, which is easy to see by the smell. Thus, the presence of any ammonium salt in a solution can be detected by heating the solution with an alkali (reaction to ammonium).

Ammonium salts are thermally unstable. When heated they decompose. This decomposition may occur reversibly or irreversibly. Ammonium salts, the anion of which is not an oxidizing agent or only weakly exhibits oxidizing properties, decompose reversibly. For example, when heated, ammonium chloride sublimes - it decomposes into ammonia and hydrogen chloride, which on the cold parts of the vessel recombine into ammonium chloride:

During the reversible decomposition of ammonium salts formed by non-volatile acids, only ammonia evaporates. However, the decomposition products - ammonia and acid - when mixed, recombine with each other. Examples include the decomposition reactions of ammonium sulfate or ammonium phosphate.

Ammonium salts, the anion of which exhibits more pronounced oxidizing properties, decompose irreversibly: a redox reaction occurs, during which ammonium is oxidized and the anion is reduced. Examples include the decomposition (§ 136) or decomposition of ammonium nitrate:

Ammonia and ammonium salts are widely used. As already mentioned, ammonia, even at low pressure, easily turns into liquid. Since a large amount of heat (1.37) is absorbed during the evaporation of liquid ammonia, liquid ammonia is used in various refrigeration devices.

Aqueous solutions of ammonia are used in chemical laboratories and industries as a weak, highly volatile base; They are also used in medicine and in everyday life. But most of the ammonia produced in industry is used for the preparation of nitric acid, as well as other nitrogen-containing substances. The most important of these include nitrogen fertilizers, primarily ammonium sulfate and nitrate and urea (p. 427).

Ammonium sulfate serves as a good fertilizer and is produced in large quantities.

Ammonium nitrate is also used as a fertilizer; The percentage of assimilable nitrogen in this salt is higher than in other nitrates or ammonium salts. In addition, ammonium nitrate forms explosive mixtures with flammable substances (ammonals) used for blasting.

Ammonium chloride, or ammonia, is used in dyeing, calico printing, soldering and tinning, as well as in galvanic cells. The use of ammonium chloride in soldering is based on the fact that it helps remove oxide films from the metal surface, so that the solder adheres well to the metal. When a highly heated metal comes into contact with ammonium chloride, the oxides located on the surface of the metal are either reduced or turn into chlorides. The latter, being more volatile than oxides, are removed from the metal surface. For the case of copper and iron, the main processes occurring can be expressed by the following equations:

The first of these reactions is redox: copper, being a less active metal than iron, is reduced by ammonia, which is formed when heated.

Liquid ammonia and solutions of ammonium salts saturated with it are used as fertilizers. One of the main advantages of such fertilizers is their increased nitrogen content.

Ammonium salts– complex substances, including ammonium cations NH4+ and acid residues.

Physical properties: ammonium salts are crystalline solids that are highly soluble in water.

Chemical properties: Ammonium has the properties of a metal, therefore the structure of its salts is similar to alkali metal salts, since NH4+ ions and alkali metal (potassium) ions have approximately the same radii. Ammonium does not exist in its free form, because it is chemically unstable and instantly decomposes into ammonia and hydrogen. Evidence of the metallic nature of ammonium is the presence of ammonium amalgam - an alloy of ammonium with mercury, similar to that of alkali metals. When treating ammonium amalgam with a cold solution of copper sulfate, the amalgam will displace nth quantity copper:

Ammonium salts have an ionic lattice and have all the properties of typical salts:

1) are strong electrolytes - they undergo dissociation in aqueous solutions, forming an ammonium cation and an acid anion:

2) undergo hydrolysis (salt of a weak base and a strong acid):

acidic environment, pH

3) enter into an exchange reaction with acids and salts:


4) interact with alkali solutions to form ammonia - a qualitative reaction to ammonium ion:

ammonium salts are determined by the smell of ammonia released as a result of the reaction, as well as by the blue color of litmus;

5) decompose when heated:

Receipt: NH3 + HNO3 = NH4NO3 (ammonium nitrate); 2NH4OH + H2SO4 = (NH4)2SO4 (ammonium sulfate) + 2H2O.

Application: ammonium salts are widely used in practice: ammonium sulfate - (NH4)2SO4, ammonium nitrate - NH4NO3, ammonium dihydrogen phosphate - NH4H2PO4 and ammonium hydrogen phosphate - (NH4)2HPO4 are used as mineral fertilizer. The advantage of the fertilizer is its increased ammonia content. Ammonium chloride (NH4Cl) is used - ammonia.

IN chemical compounds NH4+ manifests itself as a positive singly charged alkali metal cation. When interacting with acidic anions, substances with a crystalline structure are formed - salts NH4N03, chloride NH4C1, sulfate (NH4)2S04, which consist of an acid anion and an ammonium group.

Ammonium salts are obtained during the reaction of ammonia with acids.

Ammonium salts, like, in principle, most alkali metals, can dissociate (split into cations and anions) in aqueous solutions into ions: ammonium salts, like alkali metal salts, dissociate into ions:

NH4N03 ↔ NH4++ NO3-

During the heating process, dry ammonium salts decompose into ammonia and acid; This process is also called thermal dissociation.

The resulting acid (for example, hydrochloric acid) evaporates together with NH3, and upon cooling combines with ammonia to form a salt. It should be said that a reversible decomposition process is also possible:

NH3 + HCl ↔ NH4Cl

Thus, when heated, ammonia sublimes, however, after some time, a white coating of ammonium chloride appears on the upper parts of the test tube. As a result of the formation of a salt by a non-volatile acid, for example, (NH4)2S04, during heating only NH3 evaporates, and the acid remains. This process is called irreversible decomposition.

All ammonium salts decompose well with alkalis (when heated) releasing ammonia:

NH4+ + OH -↔NH3+ H20

This reaction is used to recognize mineral fertilizers with the ammonium form of nitrogen.

When ammonium carbonate interacts with minerals, carbon dioxide is released. When ammonium sulfate or chloride reacts with AgNO3 or BaCl, characteristic white precipitates are formed.

Ammonium salts: examples

Ammonium salts are widely used in agriculture. This is an excellent material for mineral feeding of plants (for example, ammonium sulfate - (NH4)2S04). As is known, plants can only absorb nitrogen in bound form (NO3, NH4). Therefore, nitrogen compounds are very effective. Ammonium nitrate, which contains ammonium NH4 and nitrate NO3-nitrogen, is of great importance.

Dihydrogen phosphate and ammonium hydrogen phosphate, known under the names ammophos NH4H2P04 and diammophos (NH4)2НР04, contain two chemical element plant nutrition - nitrogen and phosphorus. Ammonium salts are part of fertilizer mixtures.

When ammonia reacts with carbon (IV) oxide, urea or carbamide NH2-CO-NH2 is synthesized.

Ammonium chloride (NH4C1) or ammonia is used in galvanic cells, dyeing and calico printing, tinning and soldering. When in contact with heated metal, ammonia decomposes into hydrogen chloride and ammonia. interacts with the oxide contaminating the metal surface, resulting in the formation of a volatile salt.

Please note that solder adheres very well to the cleaned surface. NH4NO3 (ammonium nitrate) together with coal and aluminum salts are an integral part explosive- ammonal. These compounds are used in the development of rocks. (NH4НСО3) is very often used in the food and confectionery industry. This compound is also used as a preservative for different types feed NH4HCO3 has denitrifying properties, which accelerates the ripening of the crop.

Qualitative reaction to ammonium salts. When a solution of ammonium salts and alkalis is heated, it is formed which splits to release ammonia.

The presence of ammonia can be determined by its specific odor, as well as by using colored paper indicators. To carry out this reaction, you need to take 1.0 cm3 of ammonium salt solution, add 0.5 cm3 of the solution and heat it. During the heating process, ammonia is released, a characteristic odor is formed, and red litmus turns blue.