Lesson type: A lesson in transferring and acquiring new knowledge and skills.

Goals: Repeat and consolidate knowledge about the general chemical properties of acids; study the structure of the nitric acid molecule, the physical and specific chemical properties of nitric acid - its interaction with metals; introduce students to industrial and laboratory methods for producing pure nitric acid.

As a result of the lesson you need to know:

1. Composition and structure of the nitric acid molecule; the number of covalent bonds formed by a nitrogen atom and the degree of oxidation of nitrogen in a nitric acid molecule.
2. General chemical properties of nitric acid: interaction with indicators (litmus and methyl orange), with basic and amphoteric oxides, bases, with salts of weaker and more volatile acids.
3. Specific chemical properties of nitric acid: its interaction with metals.
4. Laboratory and industrial methods for producing nitric acid.

You must be able to:

1. Draw up equations of chemical reactions from the perspective of the theory of electrolytic dissociation.
2. Draw up reaction equations for the interaction of concentrated and dilute acids with metals using the electron balance method.

Methods and methodological techniques:

1. Conversation.
2. Independent work of students on drawing up equations of chemical reactions of nitric acid with metals.
3. Laboratory work to study the general chemical properties of nitric acid;
4. Drawing up a supporting outline.
5. Creative work: student’s report on obtaining nitric acid.
6. Demonstration of experiments: interaction of dilute and concentrated nitric acid with copper.
7. Show slides using a multimedia projector.
8. Mutual verification and mutual assessment of the results of independent work.

Equipment and reagents:

On students' desks: solutions of nitric acid HNO 3 (20 - 25%), indicators litmus and methyl orange, sodium hydroxide solution NaOH, copper (II) sulfate solution CuSO 4, iron (II) sulfate solution FeSO 4, copper (II) oxide CuO, aluminum oxide Al2O 3, sodium carbonate solution Na 2 CO 3, test tubes, test tube holders.
On the teacher's desk: concentrated nitric acid HNO 3 (60 - 65%), dilute nitric acid HNO 3 (30%), rack with test tubes, copper wire (pieces), gas outlet tube, crystallizer with water, test tube holder, multimedia installation (computer, projector, screen) .

Lesson plan:
The lesson plan is written on the board and printed for compiling a reference note on the students’ desks (Appendix 1)

During the classes:

I Repetition.

Teacher: In previous lessons we studied some nitrogen compounds. Let's remember them.
Student: These are ammonia, ammonium salts, nitrogen oxides.
Teacher: Which nitrogen oxides are acidic?
Student: Nitric oxide (III) N 2 O 3 - nitrous anhydride and nitrogen oxide (V) N 2 O 5 - nitric anhydride, it corresponds to nitric acid HNO3.
Teacher: What is the qualitative and quantitative composition of nitric acid?

The teacher writes the formula of nitric acid on the board and asks the student to arrange the oxidation states

Student: The molecule consists of three chemical elements: H, N, O - one hydrogen atom, one nitrogen atom and three oxygen atoms.

II Composition and structure of HNO 3

Teacher: How is a nitric acid molecule formed?

The teacher shows a presentation about nitric acid (Appendix 2 - presentation, Appendix 3 - text of explanation for the presentation)

III Physical properties:

Teacher: Now we move on to studying the physical properties of nitric acid.

Students write a short description of the physical properties of nitric acid.

Teacher on demonstration table showing what concentrated nitric acid isHNO (60 – 65%) is a colorless liquid, “smoking in air”, with a pungent odor. Concentrated 100%HNO 3 is sometimes yellowish in color because It is volatile and unstable, and at room temperature decomposes, releasing nitric oxide (IV) or “brown” gas, which is why it is stored in dark glass bottles.

The teacher writes on the blackboard the equation of the chemical reaction of the decomposition of nitric acid:

Teacher: Nitric acid is hygroscopic and miscible with water in any ratio. In aqueous solutions it is a strong electrolyte; it hardens at a temperature of – 41.6 0 C. In practice, 65% nitric acid is used, it does not smoke, unlike 100% - oh.

IV Chemical properties

Teacher: Let's move on to the next stage of the lesson. Nitric acid is a strong electrolyte. Consequently, it will have all the general properties of acids. What substances do acids react with?
Student: With indicators, with basic and amphoteric oxides, with bases, with salts of weaker and volatile acids, with metals.
Teacher: Here are the general properties of acids.

The multimedia installation is turned on. The teacher shows a presentation about the general chemical properties of acids (Appendix 4).

Teacher: Let's conduct an experimental stage of the lesson. Your task is to carry out chemical reactions that confirm the chemical properties of acids, using nitric acid as an example. You will work in groups of 4 people. On the desks are instructions for laboratory experiments (Appendix 5). In your notebooks you need to create equations for chemical reactions in molecular and ionic form.

Teacher: Let's move on to the specific chemical properties of nitric acid. It should be noted that nitric acid, both diluted and concentrated, does not release hydrogen when interacting with metals, but can release various nitrogen compounds - from ammonia to nitric oxide (IV).

The multimedia installation is turned on. The teacher shows a presentation about possible products of the reduction of nitric acid (Appendix 6).

Teacher: Let's look at the diagram. On everyone’s desk there are schemes for the reduction of nitric acid (diluted and concentrated) with metals (Appendix 7).

1. Reaction of dilute nitric acid with copper. Collection of nitric oxide (II) over water.
2. Reaction of concentrated nitric acid with copper. Obtaining nitric oxide (IV).

Write down reaction equations on the board:

Teacher: Based on the experiments, we can draw conclusions:

Teacher: Using the schemes for the reduction of concentrated and dilute nitric acid with metals, as well as the textbook on page 127, let’s move on to independent work on the options (Appendix 8). Everyone does their own version. You are offered cards - tasks. Working time is 5-7 minutes.

The multimedia installation is turned on. The teacher shows the correct answer options (Appendix 9). Students check whether the assignment is completed correctly.

V Preparation of nitric acid HNO 3

Student:(message) In the laboratory, nitric acid is prepared by reacting potassium or sodium nitrate with concentrated sulfuric acid with or without heating:

In industry, nitric acid is produced by the catalytic oxidation of ammonia synthesized from atmospheric nitrogen:

The student shows a diagram for the production of nitric acid (Appendix 10), and students write down the reaction equations in their notebooks.

VI Conclusion

Teacher: In today's lesson we learned about the composition and structure of nitric acid. We repeated and consolidated the general properties of acids using the example of nitric acid, consolidated our knowledge of the theory of TED, the theory of atomic structure and chemical bonding. We studied the specific properties of nitric acid, namely its interaction with metals. We learned about methods for producing nitric acid.

D/z:§ 33, ex. 4 on page 128 of the textbook;
problems: 4 – 35, 4 – 41 problem books;
learn notes.

Bibliography

1. Kuznetsova N.E., Titova I.M., Gara N.N., Zhegin A.Yu. Chemistry: textbook for 9th grade of general education institutions. – M.: Ventana – Graf, 2004.
2. Encyclopedia for children. Chemistry. – M.: Avanta, 2000.
3. Maksimenko O.O. Chemistry. A guide for applicants to universities. – M.: Eksmo, 2003.
4. Polosin V.S., Prokopenko V.G. Workshop on methods of teaching chemistry. Tutorial. – M.: Education, 1989.
5. Martynenko B.V. Chemistry: Acids and bases. – M.: Education, 2000.

Nitric acid and its properties.

Pure nitric acid HNO 3 is a colorless liquid. 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:

4HN0 3 = 4N0 2 + 0 2 + 2H 2 0.

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 strongest acids: in dilute solutions it completely disintegrates into H+ and N0_ ions.

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, is gradually oxidized into sulfuric acid, phosphorus into phosphoric acid.

Nitric acid acts on almost all metals (see section 11.3.2), turning them into nitrates, and some metals into oxides.

Concentrated HNO 3 passivates some metals.

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

4 +3 +2 +1 0 -3

N0 2 N 2 0 3 NO N 2 O N 2 NH 4 N0 3

Which of these substances is formed, i.e., how deeply nitric acid is reduced in a given case, depends on the nature of the reducing agent and on the reaction conditions, primarily on the concentration of the acid. The higher the HNO3 concentration, the less deeply it is reduced. When reacting with concentrated acid, NO2 is most often released. When dilute nitric acid reacts with low-active metals, for example, copper, NO is released. In the case of more active metals - iron, zinc - N2O 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.

Cu + HN0 3(conc.) - Cu(N0 3) 2 + N0 2 + H 2 0;

Cu + HN0 3 (diluted) -^ Cu(N0 3) 2 + N0 + H 2 O;

Mg + HN0 3 (diluted) -> Mg(N0 3) 2 + N 2 0 + n 2 0;

Zn + HN0 3 (very dilute) - Zn(N0 3) 2 + NH 4 N0 3 + H 2 0.

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

S + 6HNO 3 = H 2 S0 4 + 6N0 2 + 2H 2 0.

ZR + 5HN0 3 + 2N 2 0 = ZN 3 RO 4 + 5N0

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 acid and 3-4 volumes of concentrated hydrochloric acid is called aqua regia. 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 with the release of free chlorine and the formation of nitrogen chloroxide (1P), or nitrosyl chloride, N0C1:

HN0 3 + ZNS1 = C1 2 + 2H 2 0 + N0C1.

Nitrosyl chloride is an intermediate product of the reaction and decomposes:

2N0C1 = 2N0 + C1 2.

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:

Au + HN0 3 + ZNS1 = AuCl 3 + NO + 2H 2 0;

3Pt + 4HN0 3 + 12HC1 = 3PtCl 4 + 4N0 + 8H 2 0.

Nitric acid acts on many organic substances in such a way that one or more hydrogen atoms in the molecule of an organic compound are replaced by nitro groups - NO 2. This process is called nitration and is of great importance in organic chemistry.

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:

2KN0 3 = 2KN0 2 +O 2

Industrial production of nitric acid. Modern industrial methods for producing nitric acid are based on the catalytic oxidation of ammonia with atmospheric oxygen. When describing the properties of ammonia, it was indicated that it burns in oxygen, and the reaction products are water and free nitrogen. But in the presence of catalysts, the oxidation of ammonia with oxygen can proceed differently. If a mixture of ammonia and air is passed over the catalyst, then at 750 °C and a certain composition of the mixture, almost complete conversion of NH 3 to N0 occurs:

4NH 3 (r) + 5O 2 (g) = 4NO(r) + 6H 2 O (g), AN = -907 kJ.

The resulting NO2 easily transforms into NO2, which, with water in the presence of atmospheric oxygen, produces nitric acid.

Platinum-based alloys are used as catalysts for the oxidation of ammonia.

The nitric acid obtained by the oxidation of ammonia has a concentration not exceeding 60%. If necessary, it is concentrated.

The industry produces diluted nitric acid with concentrations of 55, 47 and 45%, and concentrated nitric acid - 98 and 97%. Concentrated acid is transported in aluminum tanks, diluted acid is transported in tanks made of acid-resistant steel.

Ticket 5

2. The role of iron in the life processes of the body.

Iron in the body. Iron is present in the bodies of all animals and in plants (on average about 0.02%); it is necessary mainly for oxygen metabolism and oxidative processes. There are organisms (so-called concentrators) capable of accumulating it in large quantities (for example, iron bacteria - up to 17-20% of iron). Almost all of the iron in animals and plants is bound to proteins. Iron deficiency causes growth retardation and chlorosis in plants associated with reduced chlorophyll formation. Excess Iron also has a harmful effect on plant development, causing, for example, sterility of rice flowers and chlorosis. In alkaline soils, Iron compounds are formed that are inaccessible for absorption by plant roots, and plants do not receive it in sufficient quantities; in acidic soils, iron turns into soluble compounds in excess quantities. When there is a deficiency or excess of assimilable iron compounds in the soil, plant diseases can be observed over large areas.

Iron enters the body of animals and humans with food (the richest sources in it are liver, meat, eggs, legumes, bread, cereals, spinach, and beets). Normally, a person receives 60-110 mg of Iron in their diet, which significantly exceeds their daily requirement. Absorption of iron received from food occurs in the upper part of the small intestines, from where it enters the blood in protein-bound form and is carried with the blood to various organs and tissues, where it is deposited in the form of an iron-protein complex - ferritin. The main depot of iron in the body is the liver and spleen. Due to ferritin, the synthesis of all iron-containing compounds of the body occurs: the respiratory pigment hemoglobin is synthesized in the bone marrow, myoglobin is synthesized in muscles, cytochromes and other iron-containing enzymes are synthesized in various tissues. Iron is released from the body mainly through the wall of the large intestines (in humans, about 6-10 mg per day) and to a small extent by the kidneys.

Introduction

You are interested in floriculture and came to the store to buy fertilizer for your flowers. While reviewing the various names and compositions, you noticed a bottle labeled “Nitrogen Fertilizer.” We read its composition: “Phosphorus, calcium, this and that... Nitric acid? What kind of animal is this?!” Usually one gets acquainted with nitric acid in such an environment. And many will then want to know more about it. Today I will try to satisfy your curiosity.

Definition

Nitric acid (formula HNO 3) is a strong monobasic acid. In an unoxidized state, it looks like in photo 1. Under normal conditions, it is a liquid, but it can be converted into a solid state of aggregation. And in it it resembles crystals having a monoclinic or rhombic lattice.

Chemical properties of nitric acid

It has the ability to mix well with water, where almost complete dissociation of this acid into ions occurs. Concentrated nitric acid is brown in color (photo). It is provided by the decomposition into nitrogen dioxide, water and oxygen, which occurs due to sunlight that falls on it. If you heat it up, the same decomposition will occur. All metals react with it, with the exception of tantalum, gold and platinoids (ruthenium, rhodium, palladium, iridium, osmium and platinum). However, its combination with hydrochloric acid can even dissolve some of them (this is the so-called “regia vodka”). Nitric acid, having any concentration, can act as an oxidizing agent. Many organic substances can spontaneously ignite when interacting with it. And some metals in this acid will be passivated. When exposed to them (as well as when reacting with oxides, carbonates and hydroxides), nitric acid forms its salts, called nitrates. The latter dissolve well in water. But nitrate ions are not hydrolyzed in it. If you heat the salts of this acid, their irreversible decomposition will occur.

Receipt

To produce nitric acid, synthetic ammonia is oxidized using platinum-rhodium catalysts to produce a mixture of nitrous gases, which are subsequently absorbed by water. It is also formed when potassium nitrate and iron sulfate are mixed and heated.

Application

Nitric acid is used to produce mineral fertilizers, explosives and some toxic substances. It is used to etch printing forms (etching boards, magnesium clichés, etc.), and also to acidify tinting solutions for photographs. Nitric acid is used to produce dyes and medicines, and it is also used to determine the presence of gold in gold alloys.

Physiological effects

Considering the degree of influence of nitric acid on the body, it is classified as hazard class 3 (moderately dangerous). Inhalation of its vapors leads to irritation of the respiratory tract. When nitric acid comes into contact with the skin, it leaves many long-healing ulcers. The areas of the skin where it gets in become a characteristic yellow color (photo). Scientifically speaking, a xanthoprotein reaction occurs. Nitrogen dioxide, which is produced when nitric acid is heated or decomposed in light, is very toxic and can cause pulmonary edema.

Conclusion

Nitric acid is beneficial to humans in both diluted and pure states. But most often it is found in substances, many of which are probably familiar to you (for example, nitroglycerin).

Nitric acid: properties and reactions,
underlying production

9th grade

When children come to a chemistry lesson, they want to learn new things and apply their knowledge; they especially like to independently obtain information and experiment. This lesson is structured so that, when studying new material, students can use previously acquired knowledge: the structure of the nitrogen atom, types of chemical bonds, electrolytic dissociation, redox reactions, safety precautions when conducting an experiment.

Goals. Review the classification and properties of nitrogen oxides, as well as the general properties of nitric acid in the light of the theory of electrolytic dissociation (ED). Introduce students to the oxidative properties of nitric acid using the example of the interaction of dilute and concentrated acid with metals. Give an idea of ​​methods for producing nitric acid and areas of its application.

Equipment. On each table in front of the students is a lesson plan, a diagram of the interaction of nitric acid with metals, a set of reagents, and tests to consolidate the studied material.

PLAN

Nitrogen oxides.

Composition and structure of the nitric acid molecule.

Physical properties of nitric acid.

Chemical properties of nitric acid.

Preparation of nitric acid.

Application of nitric acid.

Consolidation of material (test according to options).

DURING THE CLASSES

Nitrogen oxides

Teacher.Remember and write the formulas of nitrogen oxides. Which oxides are called salt-forming, which are called non-salt-forming? Why?

Students independently write down the formulas of the five nitrogen oxides, name them, remember nitrogen-containing oxygen acids and establish correspondence between oxides and acids. One of the students writes on the board (table).

Table

Comparison of nitrogen oxides, acids and salts

Demonstration experience:
interaction of nitrogen(IV) oxide with water

Teacher. In a vessel with NO 2 add a little water and shake the contents, then test the resulting solution with litmus.

What are we seeing? The solution turns red due to the formation of two acids.

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

The degree of oxidation of nitrogen in NO 2 equals +4, i.e. it is intermediate between +3 and +5, which are more stable in solution, therefore two acids correspond to nitric oxide (IV) - nitrous and nitric.

Composition and structure of the molecule

Teacher.Write down the molecular formula of nitric acid on the board, calculate its molecular mass and note the oxidation states of the elements. Write the structural and electronic formulas.

Students make up the following formulas (Fig. 1).

Rice. 1. Incorrect structural and electronic formulas of nitric acid

Teacher.According to these formulas, ten electrons rotate around nitrogen, but this cannot be, because... Nitrogen is in the second period and can only have a maximum of eight electrons in its outer layer. This contradiction is eliminated if we assume that a covalent bond is formed between the nitrogen atom and one of the oxygen atoms according to the donor-acceptor mechanism(Fig. 2).

Rice. 2. Electronic formula of nitric acid.
The electrons of the nitrogen atom are indicated by black dots

Then the structural formula of nitric acid could be depicted as follows(Fig. 3) :

Rice. 3. Structural formula of nitric acid
(donor-acceptor bond shown by arrow)

However, it has been experimentally proven that the double bond is evenly distributed between the two oxygen atoms. The oxidation state of nitrogen in nitric acid is +5, and the valence (note) is four, because there are only four common electron pairs.

Physical properties of nitric acid

Teacher.In front of you are bottles of diluted and concentrated nitric acid. Describe the physical properties you observe.

Students describe nitric acid as a liquid heavier than water, yellowish in color, with a pungent odor. Nitric acid solution is colorless and odorless.

Teacher. I will add that the boiling point of nitric acid is +83 °C, the freezing point is –41 °C, i.e. under normal conditions it is a liquid. The pungent odor and the fact that it turns yellow during storage is explained by the fact that the concentrated acid is unstable and partially decomposes when exposed to light or heating.

Chemical properties of acid

Teacher. Remember what substances acids interact with?(Students name.)

In front of you are the reagents, perform the listed reactions* and write down your observations (reactions must be written down in the light of TED).

Now let's turn to the specific properties of nitric acid.

We noted that the acid turns yellow during storage, now we will prove this with a chemical reaction:

4HNO3 = 2H2O + 4NO2 + O2.

(Students independently record the electronic balance of the reaction.)

Emitted "brown gas"(NO2) colors the acid.

This acid behaves particularly towards metals. You know that metals displace hydrogen from acid solutions, but this does not happen when interacting with nitric acid.

Look at the diagram on your desk (Fig. 4), which shows what gases are released when acids of various concentrations react with metals. (Work with the diagram.)

Rice. 4. Scheme of interaction of nitric acid with metals

Demonstration experience:
interaction of concentrated nitric acid with copper

A very effective demonstration of the reaction of nitric acid (conc.) with copper powder or finely chopped pieces of copper wire:

Students independently record the electronic balance of the reaction:

Acid production

Teacher. The lesson will be incomplete if we do not consider the issue of obtaining nitric acid.

Laboratory method: the effect of concentrated sulfuric acid on nitrates (Fig. 5).

NaNO 3 + H 2 SO 4 = NaHSO 4 + HNO 3.

In industry the acid is mainly produced by the ammonia method.

Rice. 5. To obtain nitric acid in the laboratory so far
It’s convenient to use old chemical utensils – a retort

The method of producing acid from nitrogen and oxygen at temperatures above 2000 °C (electric arc) is not particularly widespread.

In Russia, the history of the production of nitric acid is associated with the name of the chemist-technologist Ivan Ivanovich Andreev (1880–1919).

In 1915, he created the first installation for the production of acid from ammonia and implemented the developed method on a factory scale in 1917. The first plant was built in Donetsk.

This method includes several steps.

1) Preparation of ammonia-air mixture.

2) Oxidation of ammonia with air oxygen on a platinum mesh:

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

3) Further oxidation of nitric oxide (II) to nitric oxide (IV):

2NO + O 2 = 2NO 2.

4) Dissolving nitric oxide (IV) in water and producing acid:

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

If dissolution is carried out in the presence of oxygen, then all nitrogen oxide (IV) is converted into nitric acid.

5) The final stage of obtaining nitric acid is the purification of gases released into the atmosphere from nitrogen oxides. The composition of these gases: up to 98% nitrogen, 2–5% oxygen and 0.02–0.15% nitrogen oxides. (Nitrogen was initially in the air taken for ammonia oxidation.) If nitrogen oxides in these exhaust gases are more than 0.02%, then they are specially catalytically reduced to nitrogen, because even such small amounts of these oxides lead to major environmental problems.

After all that has been said, the question arises: why do we need acid?

Application of acid

Teacher.Nitric acid is used for the production of: nitrogen fertilizers, and primarily ammonium nitrate (how is it obtained?); explosives (why?); dyes; nitrates, which will be discussed in the next lesson.

Fixing the material

Frontal class survey

– Why is the oxidation state of nitrogen in nitric acid +5, and the valence is four?

– What metals does nitric acid not react with?

– You need to recognize hydrochloric and nitric acids; there are three metals on the table – copper, aluminum and iron. What will you do and why?

Test

Option 1

1. Which series of numbers corresponds to the distribution of electrons across energy levels in a nitrogen atom?

1) 2, 8, 1; 2) 2, 8, 2; 3) 2, 4; 4) 2, 5.

2. Complete the equations for practically feasible reactions:

1) HNO 3 (diluted) + Cu...;

2) Zn + HNO 3 (conc.) ... ;

3) HNO 3 + MgCO 3 ... ;

4) CuO + KNO 3 ... .

3. Indicate which equation illustrates one of the stages of the process of industrial production of nitric acid.

1) 4NH 3 + 5O 2 = 4NO + 6H 2 O;

2) 5HNO 3 + 3P + 2H 2 O = 3H 3 PO 4 + 5NO;

3) N 2 + O 2 = 2NO.

4. A negative oxidation state is manifested by nitrogen in the compound:

1) N 2 O; 2) NO; 3) NO 2; 4) Na 3 N.

5. The interaction of copper shavings with concentrated nitric acid leads to the formation of:

1) NO 2; 2) NO; 3) N 2; 4) NH 3.

Option 2

1. The value of the highest valency of nitrogen is:

1) 1; 2) 2; 3) 5; 4) 4.

2. Write down the possible interaction of concentrated nitric acid with the following metals: sodium, aluminum, zinc, iron, chromium.

3. Select the substances that are raw materials for the production of nitric acid:

1) nitrogen and hydrogen;

2) ammonia, air and water;

3) nitrates.

4. Concentrated nitric acid does not react with:

1) carbon dioxide;

2) hydrochloric acid;

3) carbon;

4) barium hydroxide.

5. When a very dilute acid reacts with magnesium, it forms:

1) NO 2; 2) NO; 3) N 2 O; 4) NH 4 NO 3.

* For example, you can invite the children to do the following laboratory experiments.

1) Add litmus to a test tube with a solution of nitric acid and gradually add sodium hydroxide solution. Write down your observations.

2) Place some chalk in a test tube and add dilute nitric acid.

3) Place some copper(II) oxide in a test tube and add dilute nitric acid. What color is the solution? Clamp the test tube in the holder and warm it. How does the color of the solution change? What does the color change mean? – Note edit.

One of the most important products used by humans is nitrate acid. The formula of the substance is HNO 3, and it also has various physical and chemical characteristics that distinguish it from other inorganic acids. In our article we will study the properties of nitric acid, get acquainted with the methods of its preparation, and also consider the scope of application of the substance in various industries, medicine and agriculture.

Features of physical properties

Nitric acid obtained in the laboratory, the structural formula of which is given below, is a colorless liquid with an unpleasant odor, heavier than water. It evaporates quickly and has a low boiling point of +83 °C. The compound is easily mixed with water in any proportions, forming solutions of varying concentrations. Moreover, nitrate acid can absorb moisture from the air, that is, it is a hygroscopic substance. The structural formula of nitric acid is ambiguous and can have two forms.

Nitrate acid does not exist in molecular form. In aqueous solutions of various concentrations, the substance has the form of the following particles: H 3 O + - hydronium ions and anions of the acid residue - NO 3 -.

Acid-base interaction

Nitric acid, which is one of the strongest acids, enters into exchange and neutralization. Thus, the compound participates in metabolic processes with basic oxides, resulting in the production of salt and water. The neutralization reaction is the basic chemical property of all acids. The products of the interaction of bases and acids will always be the corresponding salts and water:

NaOH + HNO 3 → NaNO 3 + H 2 O

Reactions with metals

In a molecule of nitric acid, the formula of which is HNO 3, nitrogen exhibits the highest oxidation state, equal to +5, so the substance has pronounced oxidizing properties. As a strong acid, it is capable of reacting with metals in the activity series of metals up to hydrogen. However, unlike other acids, it can also react with passive metal elements, for example, copper or silver. The reagents and products of the interaction are determined both by the concentration of the acid itself and by the activity of the metal.

Dilute nitric acid and its properties

If the mass fraction of HNO 3 is 0.4-0.6, then the compound exhibits all the properties of a strong acid. For example, it dissociates into hydrogen cations and anions of the acid residue. Indicators in an acidic environment, such as violet litmus, change their color to red in the presence of excess H + ions. The most important feature of reactions of nitrate acid with metals is the inability to release hydrogen, which is oxidized to water. Instead, various compounds are formed - nitrogen oxides. For example, in the process of interaction of silver with molecules of nitric acid, the formula of which is HNO 3, nitrogen monoxide, water and a salt - silver nitrate - are discovered. The degree of oxidation of nitrogen in the complex anion decreases as three electrons are added.

Nitrate acid reacts with active metal elements, such as magnesium, zinc, calcium, to form nitric oxide, the valency of which is the smallest, it is equal to 1. Salt and water are also formed:

4Mg + 10HNO3 = NH4NO3 + 4Mg(NO3)2 + 3H2O

If nitric acid, the chemical formula of which is HNO 3, is very dilute, in this case, the products of its interaction with active metals will be different. This may be ammonia, free nitrogen or nitric oxide (I). It all depends on external factors, which include the degree of metal grinding and the temperature of the reaction mixture. For example, the equation for its interaction with zinc will be as follows:

Zn + 4HNO 3 = Zn(NO 3) 2 + 2NO 2 + 2H 2 O

Concentrated HNO 3 (96-98%) acid in reactions with metals is reduced to nitrogen dioxide, and this usually does not depend on the position of the metal in the N. Beketov series. This happens in most cases when interacting with silver.

Let us remember the exception to the rule: concentrated nitric acid under normal conditions does not react with iron, aluminum and chromium, but passivates them. This means that a protective oxide film is formed on the surface of metals, preventing further contact with acid molecules. A mixture of the substance with concentrated chloride acid in a 3:1 ratio is called aqua regia. It has the ability to dissolve gold.

How nitrate acid reacts with nonmetals

The strong oxidizing properties of the substance lead to the fact that in its reactions with non-metallic elements, the latter transform into the form of the corresponding acids. For example, sulfur is oxidized to sulfate acid, boron to boric acid, and phosphorus to phosphate acid. The reaction equations below confirm this:

S 0 + 2HN V O 3 → H 2 S VI O 4 + 2N II O

Preparation of nitric acid

The most convenient laboratory method for obtaining a substance is the interaction of nitrates with concentrated It is carried out with low heating, avoiding an increase in temperature, since in this case the resulting product decomposes.

In industry, nitric acid can be produced in several ways. For example, obtained from air nitrogen and hydrogen. Acid production takes place in several stages. Intermediate products will be nitrogen oxides. First, nitrogen monoxide NO is formed, then it is oxidized by atmospheric oxygen to nitrogen dioxide. Finally, in a reaction with water and excess oxygen, dilute (40-60%) nitrate acid is produced from NO 2. If it is distilled with concentrated sulfate acid, the mass fraction of HNO 3 in the solution can be increased to 98.

The above-described method for the production of nitrate acid was first proposed by the founder of the nitrogen industry in Russia I. Andreev at the beginning of the 20th century.

Application

As we remember, the chemical formula of nitric acid is HNO 3. What feature of the chemical properties determines its use if nitrate acid is a large-scale product of chemical production? This is the high oxidizing ability of a substance. It is used in the pharmaceutical industry to obtain drugs. The substance serves as a starting material for the synthesis of explosive compounds, plastics, and dyes. Nitrate acid is used in military technology as an oxidizing agent for rocket fuel. A large volume of it is used in the production of the most important types of nitrogen fertilizers - saltpeter. They help increase the yield of the most important agricultural crops and increase the protein content in fruits and green mass.

Areas of application of nitrates

Having examined the basic properties, production and use of nitric acid, we will focus on the use of its most important compounds - salts. They are not only mineral fertilizers, some of them are of great importance in the military industry. For example, a mixture consisting of 75% potassium nitrate, 15% fine coal and 5% sulfur is called black powder. Ammonal, an explosive, is obtained from ammonium nitrate, as well as coal and aluminum powder. An interesting property of nitrate acid salts is their ability to decompose when heated.

Moreover, the reaction products will depend on which metal ion is included in the salt. If a metal element is located in the activity series to the left of magnesium, nitrites and free oxygen are found in the products. If the metal included in the nitrate is located from magnesium to copper inclusive, then when the salt is heated, nitrogen dioxide, oxygen and oxide of the metal element are formed. Salts of silver, gold or platinum at high temperatures form free metal, oxygen and nitrogen dioxide.

In our article, we found out what the chemical formula of nitric acid is in chemistry, and what features of its oxidizing properties are most important.