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
In terms of 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 prepared by reacting concentrated sulfuric acid with 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 cleaned of harmful impurities that poison 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.
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. Structural formula nitric acid is represented 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.
The electronic structure of nitric acid can be deduced 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 electrons of the nitrogen atom 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:
1. Dilute nitric acid exhibits all the properties of acids. It belongs to strong acids. IN aqueous solutions dissociates:
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 its 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 are dissolved in “aqua regia” - 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 that are 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.
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 is like concentrated hydrochloric acid, “smoke”, 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.
A characteristic property of 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 - to 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. When reacting 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 a reaction intermediate 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 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.
The electronic structure of the HNO 3 molecule is discussed in § 44.
Nitric acid is one of the most important nitrogen compounds: it is used in large quantities in the production of nitrogen fertilizers, explosives and organic dyes, and serves as an oxidizing agent in many chemical processes, used in the production of sulfuric acid using the nitrose method, 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 naturally 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 complete equations for these reactions himself.
Nitrous acid
HNO 2 is a weak monobasic acid that exists only in dilute aqueous solutions.
Salts of nitrous acid are called nitrites. Nitrites are much more stable than HNO 2, all of which are toxic.
Receipt:
1. N 2 O 3 + H 2 O = 2HNO 2
How else can you get nitrous acid? ()
What is the oxidation state of nitrous acid?
This means that the acid exhibits both oxidizing and reducing properties.
When exposed to stronger oxidizing agents, it is oxidized to HNO 3:
5HNO 2 +2HMnO 4 → 2Mn(NO 3) 2 + HNO 3 + 3H 2 O;
HNO 2 + Cl 2 + H 2 O → HNO 3 + 2HCl.
2HNO 2 + 2HI → 2NO + I 2 ↓ + 2H 2 O – reducing properties
Qualitative reaction for nitrite ion NO 2 – – interaction of nitrites with potassium iodide solution KI , acidified with dilute sulfuric acid.
How should starch iodine paper change color under the influence of free I 2?
Obtaining salts (nitrates and nitrites)
What are the methods for obtaining salts that you know? How can you get nitrates and nitrites?
1) Metal + non-metal = salt;
2) metal + acid = salt + hydrogen;
3) metal oxide + acid = salt + water;
4) metal hydroxide + acid = salt + water;
5) metal hydroxide + acid oxide= salt + water;
6) metal oxide + non-metal oxide = salt;
7) salt 1 + metal hydroxide (alkali) = salt 2 + metal hydroxide (insoluble base);
8) salt 1 + acid (strong) = salt 2 + acid (weak);
9) salt 1 + salt 2 = salt 3 + salt 4
10) salt 1 + metal (active) = salt 2 + metal (less active).
Specific method obtaining nitrates and nitrites:
disproportionation.
In the presence of excess oxygen
Salts of nitric acid - nitrates
nitrates of alkali metals, calcium, ammonium – saltpeter
KNO 3 - potassium nitrate,
NH 4 NO 3 - ammonium nitrate.
Physical properties:
All nitrates are crystalline solids, white, highly soluble in water. Poisonous!
Chemical properties of nitrates
Interaction of nitrates with metals, acids, alkalis, salts
Exercise. Note the signs of each reaction, write down the molecular and ionic equations, corresponding to the schemes:
Cu(NO 3) 2 + Zn…,
AgNO 3 + HCl…,
Cu(NO 3) 2 + NaOH…,
AgNO 3 + BaCl 2 ... .
Nitrate decomposition
When solid nitrates are heated, they all decompose with the release of oxygen (ammonium nitrate is an exception), and they can be divided into three groups.
The first group consists of alkali metal nitrates
2KNO 3 = 2KNO 2 + O 2.
The second group from alkaline earth metals up to and including copper
2Сu(NO 3) 2 = 2СuО + 4NO 2 + O 2,
The third Me group after Cu
Hg(NO 3) 2 = Hg + 2NO 2 + O 2,
Why is there a lot of nitrogen in nature (it is part of the atmosphere), and why do plants often produce poor harvests due to nitrogen starvation? (Plants cannot absorb molecular nitrogen from the air. With a lack of nitrogen, the formation of chlorophyll is delayed, the growth and development of the plant is delayed.)
Name the methods for assimilating atmospheric nitrogen.
(Part of the bound nitrogen enters the soil during thunderstorms. Legumes, on the roots of which nodule bacteria develop, capable of fixing atmospheric nitrogen, converting it into compounds available to plants.)
When harvesting crops, people annually take away huge amounts of bound nitrogen with them. He covers this loss by applying not only organic, but also mineral fertilizers (nitrate, ammonia, ammonium). Nitrogen fertilizers are applied to all crops. Nitrogen is absorbed by plants in the form of ammonium cation and nitrate anion NO 3 –.
Student reports
The effect of nitrates on environment and the human body
First aid for nitrate poisoning
Reasons for the accumulation of nitrates in vegetables and methods for growing environmentally friendly crop products
Three of the five nitrogen oxides react with water to form nitrous H1N0 2 and nitric acid HN0 3.
Nitrous acid is weak and unstable. It may be present only in small concentrations in a cooled aqueous solution. In practice, it is obtained by the action of sulfuric acid on a salt solution (most often NaN0 2) when cooled to almost 0°C. When you try to increase the concentration of nitrous acid, a blue liquid - nitric oxide (III) - is released from the solution to the bottom of the vessel. As the temperature increases, nitrous acid decomposes but the reaction
Nitric oxide (1N) reacts with water, giving two acids (see above). But taking into account the decomposition of nitrous acid, the total reaction of N 2 0 4 with water when heated is written as follows:
Salts of nitrous acid (nitrites) are quite stable. Potassium or sodium nitrites can be obtained by dissolving nitric oxide (1N) in alkali:
The formation of a mixture of salts is quite understandable, since when reacting with water, N 2 0 4 forms two acids. Neutralization with alkali prevents the decomposition of unstable nitrous acid and leads to a shift in the equilibrium of the reaction of N 2 0 4 with water completely to the right.
Alkali metal nitrites are also obtained from the thermal decomposition of their nitrates: 
Salts of nitrous acid are highly soluble in water. The solubility of some nitrites is exceptionally high. For example, at 25°C the solubility coefficient of potassium nitrite is 314, i.e. 314 g of salt dissolves in 100 g of water. Alkali metal nitrites are thermally stable and melt without decomposition.
In an acidic environment, nitrites act as fairly strong oxidizing agents. In fact, the resulting weak nitrous acid exhibits oxidizing properties. Iodine is released from iodide solutions:
Iodine is detected by its color, and nitric oxide by its characteristic odor. Nitrogen moves from CO+3 in CO +2.
Oxidizing agents stronger than nitrous acid oxidize nitrites to nitrates. In an acidic environment, a solution of potassium permanganate becomes discolored when sodium nitrite is added:
Nitrogen moves from CO+3 in CO+5. Thus, nitrous acid and nitrites exhibit redox duality.
Nitrites are poisonous because they oxidize iron (II) in hemoglobin to iron (H1) and hemoglobin loses its ability to attach and carry oxygen in the blood. The use of large amounts of nitrogen fertilizers significantly accelerates plant growth, but at the same time they contain high concentrations of nitrates and nitrites. Consumption of vegetables and berries grown in this way (watermelons, melons) leads to poisoning.
Huge practical significance has nitric acid. Its properties combine acid strength (almost complete ionization in aqueous solution), strong oxidizing properties and the ability to transfer the nitro group N0 2 + to other molecules. Nitric acid is used in large quantities to produce fertilizers. In this case, it serves as a source of nitrogen necessary for plants. It is used to dissolve metals and obtain highly soluble salts - nitrates.
An extremely important use of nitric acid is the nitration of organic substances to obtain a variety of organic products containing nitro groups. Among organic nitro compounds there are medicinal substances, dyes, solvents, explosives. Every year, global production of nitric acid exceeds 30 million tons.
In the period before the industrial development of ammonia synthesis and its oxidation, nitric acid was obtained from nitrates, for example from Chilean nitrate NaN0 3. Saltpeter was heated with concentrated sulfuric acid:
The released nitric acid vapors in the cooled receiver condense into a liquid with a high HN0 3 content.
Currently, nitric acid is produced using various variants of the method, in which the starting material is nitric oxide (N). As follows from consideration of the properties of nitrogen, its oxide NO can be obtained from nitrogen and oxygen at temperatures above 2000°C. Maintaining such a high temperature requires a lot of energy. The method was technically implemented in 1905 in Norway. The heated air passed through the voltaic arc combustion zone at a temperature of 3000-3500°C. The gases leaving the device contained only 2-3% nitrogen oxide (N). By 1925, the world production of nitrogen fertilizers using this method reached 42,000 tons. By modern scales of fertilizer production, this is very little. Subsequently, the expansion of nitric acid production followed the path of ammonia oxidation to nitrogen oxide (N).
Normal combustion of ammonia produces nitrogen and water. But when the reaction is carried out at a lower temperature using a catalyst, the oxidation of ammonia ends with the formation of NO. The appearance of NO when passing a mixture of ammonia and oxygen through a platinum mesh has been known for a long time, but this catalyst does not give a sufficiently high yield of oxide. It was possible to use this process for factory production only in the 20th century, when a more effective catalyst was found - an alloy of platinum and rhodium. The metal rhodium, which proved essential in the production of nitric acid, is approximately 10 times rarer than platinum. Reaction with a Pt/Rh catalyst in a mixture of ammonia and oxygen of a certain composition at 750°C
gives NO yield up to 98%. This process is thermodynamically less favorable than the combustion of ammonia to nitrogen and water (see above), but the catalyst ensures that the nitrogen atoms remaining after the ammonia molecule loses hydrogen quickly combine with oxygen, preventing the formation of N 2 molecules.
When a mixture containing nitric oxide (N) and oxygen is cooled, nitric oxide (N0) N0 2 is formed. Next, different options for the transformation of N0 2 are used into nitric acid. Dilute nitric acid is prepared by dissolving NQ 2 in water at elevated temperature. The reaction is given above (p. 75). Nitric acid with a mass fraction of up to 98% is obtained by reaction in a mixture of liquid N 2 0 4 with water in the presence of gaseous oxygen under high pressure. Under these conditions, nitrogen oxide (N) formed simultaneously with nitric acid has time to be oxidized by oxygen to NO 2, which immediately reacts with water. This results in the following total reaction:
The entire chain of sequential reactions of converting atmospheric nitrogen into nitric acid can be represented as follows:
The reactions of nitric oxide (NI) with water and oxygen proceed rather slowly, and it is practically impossible to achieve its complete conversion into nitric acid. Therefore, plants producing nitric acid always release nitrogen oxides into the atmosphere. Reddish smoke comes out of the factory chimney - “fox tail”. The color of the smoke is due to the presence of NO 2 . In a significant area around a large plant, forests are dying from nitrogen oxides. Coniferous tree species are especially sensitive to the effects of NO 2.
Anhydrous nitric acid is a colorless liquid with a density of 1.5 g/cm 3, boiling at 83 ° C and freezing at -41.6 ° C into a transparent crystalline substance. In air, nitric acid smokes like concentrated hydrochloric acid, since the acid vapor forms droplets of fog with water vapor in the air. Therefore, nitric acid with a low water content is called smoking. It, as a rule, has a yellow color, since it decomposes under the influence of light to form NO 2. Fuming acid is used relatively rarely.
Typically, nitric acid is produced industrially in the form of an aqueous solution with a mass fraction of 65-68%. This solution is called concentrated nitric acid. Solutions with a mass fraction of HN0 3 less than 10% - dilute nitric acid. A solution with a mass fraction of 68.4% (density 1.41 g/cm3) is azeotropic mixture, boiling at 122°C. An azeotropic mixture is characterized by the same composition of both the liquid and the vapor above it. Therefore, distillation of an azeotropic mixture does not lead to a change in its composition. In concentrated acid, along with ordinary HN0 3 molecules, there are slightly dissociated molecules of orthonitric acid H 3 N0 4.
Concentrated nitric acid passivates the surface of some metals, such as iron, aluminum, chromium. When these metals come into contact with concentrated HN() 3 chemical reaction doesn't work. This means that they stop reacting with acid. Nitric acid can be transported in steel tanks.
Both fuming and concentrated nitric acid is a strong oxidizing agent. Smoldering coal ignites when it comes into contact with nitric acid. Drops of turpentine, falling into nitric acid, ignite, forming a large flame (Fig. 20.3). Concentrated acid oxidizes sulfur and phosphorus when heated.
Rice. 20.3.
Nitric acid mixed with concentrated sulfuric acid exhibits basic properties. From the HN0 molecule 3the hydroxide ion is split off, and the nitroyl (nitronium) ion NOJ is formed:
The equilibrium concentration of nitronium is small, but such a mixture nitrates organic substances with the participation of this ion. From this example it follows that, depending on the nature of the solvent, the behavior of the substance can radically change. In water HN0 3 exhibits the properties of a strong acid, and in sulfuric acid it turns out to be a base.
In dilute aqueous solutions, nitric acid is almost completely ionized.
In concentrated solutions of nitric acid, HN0 3 molecules act as an oxidizing agent, and in dilute solutions, N0 3 ions act as an oxidizing agent, supported by an acidic environment. Therefore, nitrogen is reduced to different products depending on the acid concentration and the nature of the metal. In a neutral environment, i.e., in salts of nitric acid, the NO 3 ion becomes a weak oxidizing agent, but when a strong acid is added to neutral solutions of nitrates, the latter act as nitric acid. According to the strength of oxidizing properties in an acidic environment, the NO 3 ion stronger than H+. This leads to the following important corollary.
When nitric acid acts on metals, various nitrogen oxides are released instead of hydrogen, and in reactions with active metals, nitrogen is reduced to the NH* ion.
Let us consider the most important examples of reactions of metals with nitric acid. Copper in a reaction with dilute acid reduces nitrogen to NO (see above), and in a reaction with concentrated acid - to NO 2:
Iron is passivated with concentrated nitric acid, and with medium concentration acid it is oxidized to the oxidation state +3:
Aluminum reacts with highly dilute nitric acid without evolution of gas, since the nitrogen is reduced to CO-3, forming ammonium salt:
Salts of nitric acid, or nitrates, are known for all metals. The old name for some nitrates is often used - saltpeter(sodium nitrate, potassium nitrate). This is the only family of salts in which all salts are soluble in water. The N0 3 ion is not colored. Therefore, nitrates either turn out to be colorless salts or have the color of the cation included in their composition. Most nitrates are isolated from aqueous solutions in the form of crystalline hydrates. Anhydrous nitrates are NH 4 N0 3And alkali metal nitrates, except LiN0 3*3H 2 0.
Nitrates are often used to carry out exchange reactions in solutions. Alkali metal, calcium and ammonium nitrates are used in large quantities as fertilizers. For several centuries, potassium nitrate was of great importance in military affairs, as it was a component of the only explosive composition - gunpowder. It was obtained mainly from horse urine. The nitrogen contained in the urine, with the participation of bacteria in special saltpeter heaps, turned into nitrates. When the resulting liquid was evaporated, potassium nitrate crystallized first. This
The example shows how limited the sources of nitrogen compounds were before the industrial development of ammonia synthesis.
Thermal decomposition of nitrates occurs at temperatures below 500°C. When nitrates of active metals are heated, they transform into nitrites with the release of oxygen (see above). Nitrates of less active metals upon thermal decomposition give metal oxide, nitric oxide (1 U) and oxygen:
HNO3, an oxygen-containing monobasic strong acid. Solid nitric acid forms two crystal modifications with monoclinic and orthorhombic lattices.
Nitric acid mixes with water in any ratio. In aqueous solutions, it almost completely dissociates into ions.
It is obtained by the catalytic oxidation of synthetic ammonia on platinum-rhodium catalysts (Haber method) to a mixture of nitrogen oxides (nitrous gases), with their further absorption by water
4NH3 + 5O2 (Pt) > 4NO + 6H2O
2NO + O2 > 2NO2 4NO2 + O2 + 2H2O > 4HNO3 The concentration of nitric acid obtained by this method varies, depending on the technological design of the process, from 45 to 58%. Alchemists were the first to obtain nitric acid by heating a mixture of saltpeter and iron sulfate:
4KNO3 + 2(FeSO4 7H2O) (t°) > Fe2O3 + 2K2SO4 + 2HNO3^ + NO2^ + 13H2O
Pure nitric acid was first obtained by Johann Rudolf Glauber by treating nitrate with concentrated sulfuric acid:
KNO3 + H2SO4(conc.) (t°) > KHSO4 + HNO3^
By further distillation the so-called “fuming nitric acid”, containing virtually no water.
Application:
in the production of mineral fertilizers;
in the military industry;
in photography - acidification of some tinting solutions;
in easel graphics - for etching printed forms(etching boards, zincographic printing forms and magnesium clichés).
1. Dilute nitric acid exhibits all the properties of strong acids; in aqueous solutions it dissociates according to the following scheme:
HNO3 H+ + NO3–,
anhydrous acid:
2HNO3® NO2+ + NO3–+ H2O.
Gradually, especially in the light or when heated, nitric acid decomposes; during storage, the solution becomes brownish due to nitrogen dioxide:
4HNO3 4NO2 + 2H2O + O2.
2. Nitric acid reacts with almost all metals. Dilute nitric acid with alkali and alkaline earth metals, as well as with iron and zinc, forms the corresponding nitrates, ammonium nitrate or nitrogen hemicoxide, depending on the activity of the metal, and water:
4Mg + 10HNO3® 4Mg(NO3)2 + N2O + 5H2O,
With heavy metals, dilute acid forms the corresponding nitrates, water and nitrogen oxide is released, and in the case of stronger dilution, nitrogen:
5Fe + 12HNO3(ultra dil.)®5Fe(NO3)3 + N2+ 6H2O,
3Cu + 8HNO3® 3Cu(NO3)2 + 2NO + 4H2O.
Concentrated nitric acid, when interacting with alkali and alkali metals, forms the corresponding nitrates, water and nitrogen hemioxide is released:
8Na + 10HNO3® 8NaNO3 + N2O + 5H2O.
Metals such as iron, chromium, aluminum, gold, platinum, iridium, tantalum are passivated by concentrated acid, i.e. A film of oxides is formed on the surface of the metal, impermeable to acid. Other heavy metals when interacting with concentrated nitric acid, they form the corresponding nitrates, water and release nitrogen oxide or dioxide:
3Hg + 8HNO3(cold)®3Hg(NO3)2 + 2NO + 4H2O,
Hg + 4HNO3(hor.)®Hg(NO3)2 + 2NO2+ 2H2O,
Ag + 2HNO3® AgNO3 + NO2+ 2H2O.
3. Nitric acid is capable of dissolving gold, platinum and other noble metals, but in a mixture with hydrochloric acid. Their mixture in the ratio of three volumes of concentrated hydrochloric acid and one volume of concentrated nitric acid is called “aqua regia”. The effect of aqua regia is that nitric acid oxidizes hydrochloric acid to free chlorine, which combines with metals:
HNO3 + HCl ® Cl2 + 2H2O + NOCl,
2NOCl ® 2NO + Cl2.
Aqua regia is capable of dissolving gold, platinum, rhodium, iridium and tantalum, which do not dissolve in nitric, much less hydrochloric acid:
Au + HNO3 + 3HCl ® AuCl3 + NO + 2H2O,
HCl + AuCl3® H;
3Pt + 4HNO3 + 12HCl ® 3PtCl4 + 4NO + 8H2O,
2HCl + PtCl4® H2.
4.Non-metals are also oxidized with nitric acid to the corresponding acids; dilute acid releases nitric oxide:
3P + 5HNO3 + 2H2O ® 3H3PO4 + 5NO,
concentrated acid releases nitrogen dioxide:
S + 6HNO3® H2SO4 + 6NO2+ 2H2O,
nitric acid is also capable of oxidizing some inorganic compounds:
3H2S + 8HNO3® 3H2SO4 + 8NO + 4H2O.
HNO2 is a weak monoprotic 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. Nitrates are much more stable than HNO2, all of which are toxic.
In the gas phase, the planar nitrous acid molecule exists in two configurations, cis and trans. At room temperature the trans isomer predominates
Chem. saints
In aqueous solutions there is an equilibrium:
2HNO2 - N2O3 + H2O - NO^ + NO2^ + H2O
When the solution is heated, nitrous acid decomposes, releasing NO and NO2:
3HNO2 - HNO3 + 2NO^ + H2O.
HNO2 is a little stronger acetic acid. Easily displaced by more strong acids from salts:
H2SO4 + Ba(NO2)2 > BaSO4v + HNO2.
Nitrous acid exhibits both oxidizing and reducing properties. Under the influence of stronger oxidizing agents (H2O2, KMnO4) it is oxidized to HNO3:
2HNO2 + 2HI > 2NO^ + I2v + 2H2O;
5HNO2 + 2HMnO4 >2Mn(NO3)2 + HNO3 + 3H2O;
HNO2 + Cl2 + H2O > HNO3 + 2HCl.
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.
Receipt:
N2O3 + H2O 2HNO2,
NaNO2 + H2SO4 (0° C)® NaHSO4 + HNO2
AgNO2 + HCl ® AgCl + HNO2
Properties of salts
All nitrates are highly soluble in water. With increasing temperature, their solubility increases greatly. When heated, nitrates decompose, releasing oxygen. Nitrates of ammonium, alkali and alkaline earth metals are called nitrates, for example NaNO3 - sodium nitrate (Chilean nitrate), KNO3 - potassium nitrate, NH4NO3 - ammonium nitrate. Nitrates are produced by the action of nitric acid HNO3 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.
Alkali metal nitrates decompose to nitrites with the release of oxygen (and with prolonged heating they decompose stepwise into metal oxide, molecular nitrogen and oxygen, which is why they are good oxidizing agents).
Medium activity metal nitrates decompose when heated to metal oxides, releasing nitrogen dioxide and oxygen.
Nitrates of the least active metals (noble metals) decompose mainly to free metals with the release of nitrogen dioxide and oxygen.
Nitrates are quite strong oxidizing agents in solid state(usually in the form of a melt), but have practically no oxidizing properties in solution, unlike nitric acid.
Nitrite is a salt of nitrous acid HNO2. Nitrites are thermally less stable than nitrates. They are used in the production of azo dyes and in medicine.