“Salt of nitric acid” - What happens when sodium nitrate decomposes? Physical properties of nitrates. Specify the oxidizing agent and reducing agent. Chemical properties of nitrates. Know and be able to. A solution of nitric acid reacts with each of the substances. What conclusions did the young chemist come to? All nitrates are thermally unstable. Interesting story.

“Theories of acids and bases” - For example, which acid is stronger than acetic acid (CH3COOH or chloroacetic ClCH2COOH? 2. Addition reactions. The strength of the base R3N in water can be assessed by considering the equilibrium: Gilbert Newton Lewis. A measure of acidity is an equilibrium constant called the acidity constant (Ka) . Svante August Arrhenius.

“Acetic acid” - What are acids? Not all fruits and vegetables contain acids. Acids. Formic acid was first obtained in its pure form in 1749. Andreas Sigismund Marggraff. Journey into the world of acids. How to detect acids? Formic acid solution. History of the discovery of acids. Acids have a similar composition. What substances are acids?

“Acid rain” - New ozone holes are expected to appear in the Baltic region. Aluminum can cause disease. One can imagine what happens to wild animal species when forests die. We, and almost all living things, need fresh water. Along with the death of lakes, forest degradation also becomes apparent.

“Carboxylic acids” - Repeat the definition of carboxylic acids. Preparation of carboxylic acids. The structure of the carboxyl group. Carboxylic acids. What are carboxylic acids called? Nomenclature of esters. All carboxylic acids have a functional group. Carboxylic acid molecules form dimers. Chemical properties of carboxylic acids.

“Nucleic acids” - 1892 - chemist Lilienfeld isolated thymonucleic acid from the thymus gland in 1953. Nitrogenous base. Laboratory workshop. Structure of nucleotides (differences). 1868 - German chemist F. Miescher discovered nucleic acids. Biological role of nucleic acids. Comparative characteristics. The length of DNA molecules (American biologist G. Taylor).

Nitric acid.

Completed by: teacher of chemistry and biology

Muravyova Nina Ivanovna

• Nitrogen oxides
• The structure of the nitric acid molecule.
• Preparation of nitric acid
• Physical properties.
• Properties of nitrates.
• Laboratory experiment
• Application of nitric acid and nitrates

Nitrogen oxides

Table

Comparison of nitrogen oxides, acids and salts

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

The structure of the nitric acid molecule.

Structural formula of nitric acid

Preparation of nitric acid

In the laboratory NaNO 3 (TV.) + H 2 SO 4 (END) → NaHSO 4 + HNO 3

In industry

4NH 3 + 5O 2 →4NO + 6H 2 O + Q

2NO + O 2 → 2NO 2 (when cooling)

4NO 2 + O 2 + 2H 2 O ↔ 4HNO 3 + Q

Preparation of nitric acid by oxidation of ammonia with atmospheric oxygen.

Ammonia-air mixture

Scheme for producing nitric acid in industry

2 NO2+O2 →2NO2

3NO2+H2O →2HNO 3 +NO

catalyst

Contact device

Oxidation tower

Absorption tower

Contact device

Ammonia-air

Catalyst

Nitrous gases

Physical properties

Pure nitric acid is a colorless, fuming liquid with a strong, irritating odor. Concentrated nitric acid is usually yellow in color. This color is given to it by nitric oxide (IV), which is formed due to the partial decomposition of nitric acid and dissolves in it.

• Nitric acid is a strong oxidizing agent, concentrated nitric acid oxidizes sulfur to sulfuric acid, and phosphorus to phosphoric acid, some organic compounds (for example, amines and hydrazines, turpentine) self-ignite upon contact with concentrated nitric acid.

Properties of nitrates

Me is to the left of Mg

MeNO 2 + O 2 ↓

Me is between Mg and Cu

MeO + NO 2 + O 2

Me are to the right of Cu

Me + NO 2 + O 2

• Carefully add several thin pieces of copper wire to a test tube containing concentrated nitric acid. The reaction occurs without heating, students observe a change in the color of the solution and the release of red-brown gas NO2

check yourself

Cu + HNO 3 (END) = Cu(NO 3 ) 2 + NO 2 +H 2 O

• Carefully add several thin pieces of copper wire to a test tube containing dilute nitric acid. The reaction occurs when heated. Observe the color change of the solution and the release of colorless NO gas
• Write an equation for the reaction that occurs

Test yourself

Cu + HNO3(detailed) = Cu(NO3)2 + NO + H2O

Cu 0 – 2e = Cu +1 3 the reducing agent is oxidized

N +5 + 3e = N +2 2 the oxidizing agent is reduced

3Cu + 8HNO3 = 3Cu(NO3)2 + 2NO + 4H2O

Application of nitric acid and nitrates

MEDICINES

DYES

COLLODION

EXPLOSIVES

PHOTO FILM

AQUA REGIA

MINERAL FERTILIZERS

• 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?

Physical and physicochemical properties The molecule has a flat structure (bond lengths in nm): nitrogen in nitric acid is tetravalent, oxidation state +5. nitric acid is a colorless liquid that fumes in air, concentrated nitric acid is usually yellow in color (highly concentrated HNO3 is usually brown in color due to the decomposition process occurring in the light: 4HNO3 == 4NO2  + 2H2O + O2  ) melting point -41.59°C, boiling point +82.6°C with partial decomposition. The solubility of nitric acid in water is unlimited. In aqueous solutions, it almost completely dissociates into ions. Forms an azeotropic mixture with water.

Chemical properties When heated, nitric acid decomposes according to the same reaction.

Nitric acid in any concentration exhibits the properties of an oxidizing acid, with nitrogen being reduced to an oxidation state from +4 to -3. The depth of reduction depends primarily on the nature of the reducing agent and the concentration of nitric acid. Nitric acid in any concentration exhibits the properties of an oxidizing acid, with nitrogen being reduced to an oxidation state from +4 to -3. The depth of reduction depends primarily on the nature of the reducing agent and the concentration of nitric acid. As an oxidizing acid, HNO3 interacts: a) with metals in the voltage series to the right of hydrogen: Concentrated HNO3 Cu + 4HNO3(60%) = Cu(NO3)2 + 2NO2  + 2H2O Dilute HNO3 3Cu + 8HNO3(30%) = 3Cu(NO3)2 + 2NO  + 4H2O b) with metals in the voltage series to the left of hydrogen: Zn + 4HNO3(60%) = Zn(NO3)2 + 2NO2  + 2H2O 3Zn + 8HNO3(30%) = 3Zn(NO3)2 + 2NO  + 4H2O 4Zn + 10HNO3(20%) = 4Zn(NO3) 2 + N2O  + 5H2O 5Zn + 12HNO3 = 5Zn(NO3) 2 + N2  + 6H2O d 4Zn + 10HNO3(3%) = 4Zn(NO3)2 + NH4NO3 + 3H2O All the above equations reflect only the dominant course of the reaction. This means that under given conditions there are more products of this reaction than products of other reactions, for example, when zinc reacts with nitric acid (mass fraction of nitric acid in solution 0.3), the products will contain the most NO, but will also contain (only in smaller quantities) and NO2, N2O, N2 and NH4NO3.

Nitrate HNO3 is a strong acid. Its salts - nitrates - are obtained by the action of HNO3 on metals, oxides, hydroxides or carbonates. All nitrates are highly soluble in water. Salts of nitric acid - nitrates - decompose irreversibly when heated, the decomposition products are determined by the cation: a) nitrates of metals located in the voltage series to the left of magnesium: 2NaNO3 = 2NaNO2 + O2 b) nitrates of metals located in the voltage series between magnesium and copper: 4Al(NO3 )3 = 2Al2O3 + 12NO2 + 3O2 c) nitrates of metals located in the voltage series to the right of mercury: 2AgNO3 = 2Ag + 2NO2 + O2 d) ammonium nitrate: NH4NO3 = N2O + 2H2O Nitrates in aqueous solutions practically do not exhibit oxidizing properties, but at high temperature in the solid state, nitrates are strong oxidizing agents, for example: Fe + 3KNO3 + 2KOH = K2FeO4 + 3KNO2 + H2O - when fusing solids.

Salts of nitric acid - nitrates - are widely used as fertilizers. Moreover, almost all nitrates are highly soluble in water, so there are extremely few of them in nature in the form of minerals; the exceptions are Chilean (sodium) nitrate and Indian nitrate (potassium nitrate). Most nitrates are obtained artificially. Glass and fluoroplastic-4 do not react with nitric acid.

Nitric acid production Industrial production. The modern method of its production is based on the catalytic oxidation of synthetic ammonia on platinum-rhodium catalysts to a mixture of nitrogen oxides, with their further absorption by water. The industrial method for producing HNO3 consists of the following main stages: 1. oxidation of ammonia to NO in the presence of a platinum-rhodium catalyst: 4NH3 + 5O2 = 4NO + 6H2O 2. oxidation of NO to NO2 in the cold under pressure (10 at, 1 MPa): 2NO + O2 = 2NO2 3. absorption of NO2 by water in the presence of oxygen: 4NO2 + 2H2O + O2= 4HNO3 Mass fraction of HNO3 in the resulting solution is about 0.6. The rarely used arc method for producing nitric acid differs only in the first stage, which consists of passing air through the flame of an electric arc: N2 + O2 = 2NO

For the first time, alchemists obtained nitric acid by heating a mixture of saltpeter and iron sulfate: For the first time, alchemists obtained nitric acid by heating a mixture of saltpeter and iron sulfate: 4KNO3 + 2(FeSO4 7H2O) (t°) → Fe2O3 + 2K2SO4 + 2HNO3 + NO2 + 13H2O Pure nitric acid the acid was first obtained by Johann Rudolf Glauber, acting on nitrate with concentrated sulfuric acid: KNO3 + H2SO4 (conc.) (t°) → KHSO4 + HNO3 Further distillation can produce the so-called. “fuming nitric acid”, containing virtually no water.

- This substance was described by the Arab chemist in the 8th century Jabir ibn Hayyan (Geber) in his work “The Coachman of Wisdom”, and since the 15th century this substance has been extracted for industrial purposes. - Thanks to this substance, Russian scientist V.F. Petrushevsky first received dynamite in 1866. - This substance is the progenitor of most explosives (for example, TNT, or tola). - This substance is a component of rocket fuel; it was used for the engine of the world's first Soviet jet aircraft, BI-1. - This substance, mixed with hydrochloric acid, dissolves platinum and gold, recognized as the “king” of metals. The mixture itself, consisting of 1 volume of this substance and 3 volumes of hydrochloric acid, is called “aqua regia”.

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Slide captions:

Preparation of nitric acid PREPARED BY: 9th grade student of gymnasium No. 1 named after. Yu.A. Gagarina Mikhalchenko Ksenia.

Physical properties of nitric acid Physical state: liquid Color: colorless Odor: pungent Density: 1.5 2 g/cm 3 Unlimited soluble in water Boiling: +82.6 °C with partial decomposition; Melting: −41.59 °C

Chemical properties of nitric acid HNO 3 is a strong monobasic acid. Highly concentrated HNO 3 is usually brown in color due to the decomposition process occurring in the light 4 HNO 3 4NO 2 + 2 H 2 O + O 2 When heated, nitric acid decomposes according to the same reaction. Nitric acid can be distilled (without decomposition) only under reduced pressure. Nitric acid in any concentration exhibits the properties of an oxidizing acid.

The most important compounds A mixture of three volumes of hydrochloric acid and one volume of nitric acid is called “Royal Vodka”. Aqua regia dissolves most metals, including gold and platinum. Its strong oxidizing abilities are due to the resulting atomic chlorine and nitrosyl chloride: Nitrates are salts of nitric acid. Nitrates are produced by the action of nitric acid HNO 3 on metals, oxides, hydroxides, and salts. Almost all nitrates are highly soluble in water. Nitrates are stable at ordinary temperatures. They usually melt at relatively low temperatures (200-600 °C), often with decomposition.

Occurrence in nature It is not found in nature in a free state, but always only in the form of nitrate salts. So, in the form of ammonium nitrate in the air and rainwater, especially after thunderstorms, then in the form of sodium nitrate in Chilean or Peruvian saltpeter and potassium and calcium nitrate in the upper layers of arable land, on the walls of stables, in the lowlands of the Ganges and other rivers of India. * Saltpeter is a trivial name for minerals containing nitrates of alkali and alkaline earth metals.

Virtual experiment Attention! Nitric acid and its vapors are very harmful, so you should work with it very carefully.

Production of nitric acid A distinction is made between the production of weak (diluted) nitric acid and the production of concentrated nitric acid. The process of producing dilute nitric acid consists of three stages: 1) conversion of ammonia to produce nitrogen oxide 4NH 3 + 5O 2 → 4NO + 6H 2 O 2) oxidation of nitric oxide to nitrogen dioxide 2NO + O 2 → 2NO 2 3) absorption of nitrogen oxides water 4NO 2 + O 2 + 2H 2 O → 4HNO 3 The total reaction of the formation of nitric acid is expressed as NH 3 + 2O 2 → HNO 3 + H 2 O

The use of nitric acid to produce: nitrogen fertilizers; Medicines Dyes Explosives Plastic masses Artificial fibers “Fuming” nitric acid is used in rocket technology as an oxidizer for rocket fuel extremely rarely in photography - diluted - acidification of some tinting solutions; in easel graphics - for etching printing forms (etching boards, zincographic printing forms and magnesium clichés). in jewelry - the main way to determine gold in a gold alloy;

On the topic: methodological developments, presentations and notes

appendix to the lesson “Nitric acid: molecular composition, physical and chemical properties.” “Nitric acid: molecular composition, physical and chemical properties.” Appendix to the lesson "Nitric acid:

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