Which compound has the oxidation state of nitrogen equal to 3. The oxidation state of nitrogen - learning to understand

There are chemical elements that exhibit different oxidation states, which allows the formation of chemical reactions a large number of compounds with certain properties. Knowing electronic structure atom, you can guess what substances will be formed.

The oxidation states of nitrogen can vary from -3 to +5, which indicates the variety of compounds based on it.

Element characteristic

Nitrogen belongs to the chemical elements located in the 15th group, in the second period in the periodic system of Mendeleev D.I. It was assigned the serial number 7 and the abbreviated letter designation N. Under normal conditions, a relatively inert element, special conditions are necessary for reactions.

It occurs naturally as a diatomic colorless gas. atmospheric air with a volume fraction of more than 75%. It is contained in the composition of protein molecules, nucleic acids and nitrogen-containing substances of inorganic origin.

Atom structure

To determine the oxidation state of nitrogen in compounds, you need to know it nuclear structure and explore electron shells.

A natural element is represented by two stable isotopes, with their mass number 14 or 15. The first nucleus contains 7 neutron and 7 proton particles, and the second contains 1 more neutron particle.

There are artificial varieties of its atom with a mass of 12-13 and 16-17, which have unstable nuclei.

When studying the electronic structure of atomic nitrogen, it can be seen that there are two electron shells (inner and outer). The 1s orbital contains one pair of electrons.

The second outer shell contains only five negatively charged particles: two in the 2s sublevel and three in the 2p orbital. The valence energy level does not have free cells, which indicates the impossibility of separating its electron pair. The 2p orbital is considered to be only half filled with electrons, which allows 3 negatively charged particles to be attached. In this case, the oxidation state of nitrogen is -3.

Considering the structure of the orbitals, we can conclude that this element with a coordination number of 4 binds to the maximum with only four other atoms. To form three bonds, an exchange mechanism is used, another one is formed in a do-nor-but-ac-chain-tor way.

Nitrogen oxidation states in different compounds

The maximum number of negative particles that its atom can attach is 3. In this case, its oxidation state is manifested equal to -3, inherent in compounds such as NH 3 or ammonia, NH 4 + or ammonium and Me 3 N 2 nitrides. The latter substances are formed with an increase in temperature by the interaction of nitrogen with metal atoms.

The largest number of negatively charged particles that an element can give off is equal to 5.

Two nitrogen atoms are able to combine with each other to form stable compounds with an oxidation state of -2. Such a bond is observed in N 2 H 4 or hydrazines, in azides of various metals or MeN 3 . The nitrogen atom adds 2 electrons to free orbitals.

There is an oxidation state of -1 when a given element receives only 1 negative particle. For example, in NH 2 OH or hydroxylamine it is negatively charged.

There are positive signs of the degree of nitrogen oxidation, when electron particles are taken from the outer energy layer. They vary from +1 to +5.

The charge 1+ is present in nitrogen in N 2 O (monovalent oxide) and in sodium hyponitrite with the formula Na 2 N 2 O 2 .

In NO (divalent oxide), the element donates two electrons and becomes positively charged (+2).

There is an oxidation state of nitrogen 3 (in the compound NaNO 2 or nitride and also in the trivalent oxide). In this case, 3 electrons are split off.

The +4 charge occurs in an oxide with valency IV or its dimer (N 2 O 4).

The positive sign of the oxidation state (+5) appears in N 2 O 5 or in pentavalent oxide, in nitric acid and its derivative salts.

Compounds from nitrogen to hydrogen

Natural substances based on the above two elements resemble organic hydrocarbons. Only hydrogen nitrogens lose their stability with an increase in the amount of atomic nitrogen.

The most significant hydrogen compounds include the molecules of ammonia, hydrazine, and hydrazoic acid. They are obtained by the interaction of hydrogen with nitrogen, and oxygen is also present in the latter substance.

What is ammonia

It is also called hydrogen nitride, and its chemical formula denoted as NH 3 with a mass of 17. Under conditions of normal temperature and pressure, ammonia has the form of a colorless gas with a pungent ammonia odor. In terms of density, it is 2 times rarer than air, it easily dissolves in the aquatic environment due to the polar structure of its molecule. Refers to low-risk substances.

In industrial volumes, ammonia is produced by catalytic synthesis from hydrogen and nitrogen molecules. There are laboratory methods for obtaining nitrite from ammonium salts and sodium.

The structure of ammonia

The pyramidal molecule contains one nitrogen and 3 hydrogen atoms. They are located in relation to each other at an angle of 107 degrees. In a tetrahedral molecule, nitrogen is located in the center. Through three unpaired p-electrons it is connected by polar bonds of a covalent nature with 3 atomic hydrogens, which each have 1 s-electron. This is how an ammonia molecule is formed. In this case, nitrogen exhibits an oxidation state of -3.

This element still has an unshared pair of electrons at the outer level, which creates a covalent bond with a hydrogen ion that has a positive charge. One element is a donor of negatively charged particles, and the other is an acceptor. This is how the ammonium ion NH 4 + is formed.

What is ammonium

It is classified as a positively charged polyatomic ion or cation. Ammonium is also classified as a chemical substance that cannot exist in the form of a molecule. It is made up of ammonia and hydrogen.

Ammonium with a positive charge in the presence of various anions with a negative sign is able to form ammonium salts, in which it behaves like metals with valence I. Also, with its participation, ammonium compounds are synthesized.

Many ammonium salts exist as crystalline, colorless substances that are readily soluble in water. If the compounds of the NH 4 + ion are formed by volatile acids, then under heating conditions they decompose with the release of gaseous substances. Their subsequent cooling leads to a reversible process.

The stability of such salts depends on the strength of the acids from which they are formed. Stable ammonium compounds correspond to a strong acid residue. For example, stable ammonium chloride is produced from hydrochloric acid. At temperatures up to 25 degrees, such salt does not decompose, which cannot be said about ammonium carbonate. The latter compound is often used in cooking for rising dough, replacing baking soda.

Confectioners simply call ammonium carbonate ammonium. This salt is used by brewers to improve the fermentation of brewer's yeast.

A qualitative reaction for the detection of ammonium ions is the action of hydroxides alkali metals to its connections. In the presence of NH 4 +, ammonia is released.

Chemical structure of ammonium

The configuration of its ion resembles a regular tetrahedron, in the center of which is nitrogen. Hydrogen atoms are located at the tops of the figure. To calculate the oxidation state of nitrogen in ammonium, you need to remember that the total charge of the cation is +1, and each hydrogen ion is missing one electron, and there are only 4 of them. The total hydrogen potential is +4. If we subtract the charge of all hydrogen ions from the charge of the cation, we get: +1 - (+4) = -3. So nitrogen has an oxidation state of -3. In this case, it adds three electrons.

What are nitrides

Nitrogen is able to combine with more electropositive atoms of metallic and non-metallic nature. As a result, compounds similar to hydrides and carbides are formed. Such nitrogen-containing substances are called nitrides. Between the metal and the nitrogen atom in the compounds, covalent, ionic and intermediate bonds are distinguished. It is this characteristic that underlies their classification.

Covalent nitrides include compounds in the chemical bond of which electrons do not transfer from atomic nitrogen, but form a common electron cloud together with negatively charged particles of other atoms.

Examples of such substances are hydrogen nitrides, such as ammonia and hydrazine molecules, as well as nitrogen halides, which include trichlorides, tribromides and trifluorides. They have a common electron pair that equally belongs to two atoms.

Ionic nitrides include compounds with chemical bond, formed by the transition of electrons from a metal element to free levels in nitrogen. Polarity is observed in the molecules of such substances. Nitrides have a nitrogen oxidation state of 3-. Accordingly, the total charge of the metal will be 3+.

Such compounds include nitrides of magnesium, lithium, zinc or copper, with the exception of alkali metals. They have a high melting point.

Intermediate nitrides include substances in which the atoms of metals and nitrogen are evenly distributed and there is no clear shift of the electron cloud. Such inert compounds include nitrides of iron, molybdenum, manganese and tungsten.

Description of trivalent nitric oxide

It is also called an anhydride derived from nitrous acid having the formula HNO 2 . Taking into account the oxidation states of nitrogen (3+) and oxygen (2-) in trioxide, the ratio of atoms of elements 2 to 3 or N 2 O 3 is obtained.

The liquid and gaseous forms of anhydride are very unstable compounds; they easily decompose into 2 different oxides with valences IV and II.

DEFINITION

Nitrogen- the seventh element Periodic table. It is located in the second period of the V group of the A subgroup. Designation - N.

Nitrogen is a typical non-metallic element, in terms of electronegativity (3.0) it is second only to fluorine and oxygen.

Natural nitrogen consists of two stable isotopes 14 N (99.635%) and 15 N (0.365%).

The nitrogen molecule is diatomic. There is a triple bond between the nitrogen atoms in the molecule, as a result of which the N 2 molecule is exceptionally strong. Molecular nitrogen is chemically inactive, weakly polarized.

IN normal conditions molecular nitrogen is a gas. The melting points (-210 o C) and boiling points (-195.8 o C) of nitrogen are very low; it is poorly soluble in water and other solvents.

The oxidation state of nitrogen in compounds

Nitrogen forms diatomic molecules of the composition N 2 due to the induction of covalent non-polar bonds, and, as is known, in compounds with non-polar bonds, the oxidation state of the elements is zero.

Nitrogen is characterized by a whole range of oxidation states, among which there are both positive and negative.

Oxidation state (-3) nitrogen manifests itself in compounds called nitrides (Mg +2 3 N -3 2, B +3 N -3), the most famous of which is ammonia (N -3 H +1 3).

Oxidation state (-2) nitrogen manifests itself in peroxide-type compounds - pernitrides, the simplest representative of which is hydrazine (diamide / hydrogen pernitride) - N -2 2 H 2.

In a compound called hydroxylamine - N -1 H 2 OH-nitrogen shows the oxidation state (-1) .

The most stable positive nitrogen oxidation states are (+3) And (+5) . He exhibits the first of them in fluoride (N +3 F -1 3), oxide (N +3 2 O -2 3), oxohalides (N +3 OCl, N +3 OBr, etc.), as well as derivatives anion NO 2 - (KN + 3 O 2, NaN + 3 O 2, etc.). The oxidation state (+5) nitrogen shows in oxide N + 5 2 O 5, oxonitride N + 5 ON, dioxofluoride N + 5 O 2 F, as well as in trioxonitrate (V) -ion NO 3 - and dinitridonitrate (V) -ion NH 2 -.

Nitrogen also exhibits oxidation states (+1) - N +1 2 O, (+2) - N +2 O and (+4) N +4 O 2 in their compounds, but much less frequently.

Examples of problem solving

EXAMPLE 1

The task	Indicate the oxidation states of oxygen in the compounds: La 2 O 3 , Cl 2 O 7 , H 2 O 2 , Na 2 O 2 , BaO 2 , KO 2 , KO 3 , O 2 , OF 2 .
Answer	Oxygen forms several types of binary compounds in which it exhibits characteristic oxidation states. So, if oxygen is part of oxides, then its oxidation state is (-2), as in La 2 O 3 and Cl 2 O 7. In peroxides, the oxidation state of oxygen is (-1): H 2 O 2 , Na 2 O 2 , BaO 2 . In combination with fluorine (OF 2), the oxidation state of oxygen is (+2). The oxidation state of an element in simple matter is always zero (O o 2). Substances of the composition KO 2 and KO 3 are superperoxide (superoxide) and potassium ozonide, in which oxygen exhibits fractional values ​​of oxidation states: (-1/2) and (-1/3).
Answer	(-2), (-2), (-1), (-1), (-1), (-1/2), (-1/3), 0 and (+2).

EXAMPLE 2

The task	Indicate the oxidation states of nitrogen in the compounds: NH 3 , N 2 H 4 , NH 2 OH, N 2 , N 2 O, NO, N 2 O 3 , NO 2 , N 2 O 5 .
Solution	The oxidation state of an element in a simple substance is always zero (N o 2). It is known that in oxides the oxidation state of oxygen is (-2). Using the electroneutrality equation, we determine that the oxidation states of nitrogen in oxides are: N +1 2 O, N +2 O, N +3 2 O 3, N +4 O 2, N +5 2 O 5.

Nitrogen is perhaps the most common chemical element in all solar system. To be more specific, nitrogen is the 4th most abundant. Nitrogen in nature is an inert gas.

This gas is colorless and odorless and very difficult to dissolve in water. However, nitrate salts tend to react very well with water. Nitrogen has a low density.

Nitrogen is an amazing element. There is an assumption that it got its name from the ancient Greek language, which means “lifeless, spoiled” in translation from it. Why such a negative attitude towards nitrogen? After all, we know that it is part of proteins, and breathing without it is almost impossible. Nitrogen plays an important role in nature. But in the atmosphere this gas is inert. If it is taken as it is in its original form, then many side effects are possible. The victim may even die from suffocation. After all, nitrogen is called lifeless because it does not support combustion or respiration.

Under normal conditions, such a gas reacts only with lithium, forming a compound such as lithium nitride Li3N. As we can see, the oxidation state of nitrogen in such a compound is -3. With other metals, and of course, it also reacts, but only when heated or when using various catalysts. By the way, -3 is the lowest oxidation state of nitrogen, since only 3 electrons are needed to completely fill the outer energy level.

This indicator has various meanings. Each oxidation state of nitrogen has its own compound. It is better to just remember such connections.

5 - highest degree nitrogen oxidation. Occurs in and in all nitrate salts.

Nitrogen- element of the 2nd period of the V A-group Periodic system, serial number 7. Electronic formula atom [ 2 He] 2s 2 2p 3, the characteristic oxidation states are 0, -3, +3 and +5, less often +2 and +4, etc. the N v state is considered relatively stable.

Nitrogen oxidation state scale:
+5 - N 2 O 5, NO 3, NaNO 3, AgNO 3

3 - N 2 O 3 , NO 2 , HNO 2 , NaNO 2 , NF 3

3 - NH 3, NH 4, NH 3 * H 2 O, NH 2 Cl, Li 3 N, Cl 3 N.

Nitrogen has a high electronegativity (3.07), the third after F and O. It exhibits typical non-metallic (acidic) properties, while forming various oxygen-containing acids, salts and binary compounds, as well as the ammonium cation NH 4 and its salts.

In nature - seventeenth by chemical abundance element (ninth among non-metals). A vital element for all organisms.

N 2

Simple substance. It consists of non-polar molecules with a very stable N≡N ˚σππ bond, which explains the chemical inertness of the element under normal conditions.

A colorless, tasteless, odorless gas that condenses to a colorless liquid (unlike O2).

The main component of air is 78.09% by volume, 75.52 by mass. Nitrogen boils out of liquid air before oxygen does. Slightly soluble in water (15.4 ml / 1 l H 2 O at 20 ˚C), the solubility of nitrogen is less than that of oxygen.

At room temperature, N 2 reacts with fluorine and, to a very small extent, with oxygen:

N 2 + 3F 2 \u003d 2NF 3, N 2 + O 2 ↔ 2NO

The reversible reaction of obtaining ammonia proceeds at a temperature of 200˚C, under pressure up to 350 atm, and always in the presence of a catalyst (Fe, F 2 O 3 , FeO, in the laboratory at Pt)

N 2 + 3H 2 ↔ 2NH 3 + 92 kJ

In accordance with the Le Chatelier principle, an increase in the yield of ammonia should occur with an increase in pressure and a decrease in temperature. However, the reaction rate at low temperatures is very low, so the process is carried out at 450-500 ˚C, reaching a 15% yield of ammonia. Unreacted N 2 and H 2 return to the reactor and thereby increase the extent of the reaction.

Nitrogen is chemically passive with respect to acids and alkalis, does not support combustion.

Receipt in industry- fractional distillation of liquid air or chemical removal of oxygen from air, for example, by the reaction 2C (coke) + O 2 \u003d 2CO when heated. In these cases, nitrogen is obtained, which also contains impurities of noble gases (mainly argon).

In the laboratory, small amounts of chemically pure nitrogen can be obtained by a switching reaction with moderate heating:

N -3 H 4 N 3 O 2 (T) \u003d N 2 0 + 2H 2 O (60-70)

NH 4 Cl(p) + KNO 2 (p) = N 2 0 + KCl + 2H 2 O (100˚C)

It is used for the synthesis of ammonia. Nitric acid and other nitrogen-containing products as an inert medium for chemical and metallurgical processes and storage of flammable substances.

NH 3

Binary compound, nitrogen oxidation state is - 3. A colorless gas with a sharp characteristic odor. The molecule has the structure of an incomplete tetrahedron [: N(H) 3 ] (sp 3 hybridization). The presence of nitrogen in the NH 3 molecule of a donor pair of electrons in the sp 3 hybrid orbital causes characteristic reaction addition of a hydrogen cation to form a cation ammonium NH4. It liquefies under positive pressure at room temperature. In the liquid state, it is associated by hydrogen bonds. Thermally unstable. Let's well dissolve in water (more than 700 l/1 l of H 2 O at 20˚C); the proportion in the saturated solution is 34% by weight and 99% by volume, pH= 11.8.

Very reactive, prone to addition reactions. Burns in oxygen, reacts with acids. Shows reducing (due to N -3) and oxidizing (due to H +1) properties. It is dried only with calcium oxide.

Qualitative reactions - the formation of white "smoke" upon contact with gaseous HCl, blackening of a piece of paper moistened with a solution of Hg 2 (NO3) 2.

An intermediate product in the synthesis of HNO 3 and ammonium salts. It is used in the production of soda, nitrogen fertilizers, dyes, explosives; liquid ammonia is a refrigerant. Poisonous.
Equations of the most important reactions:

2NH 3 (g) ↔ N 2 + 3H 2
NH 3 (g) + H 2 O ↔ NH 3 * H 2 O (p) ↔ NH 4 + + OH -
NH 3 (g) + HCl (g) ↔ NH 4 Cl (g) white "smoke"
4NH 3 + 3O 2 (air) = 2N 2 + 6 H 2 O (combustion)
4NH 3 + 5O 2 = 4NO+ 6 H 2 O (800˚C, cat. Pt/Rh)
2 NH 3 + 3CuO = 3Cu + N 2 + 3 H 2 O (500˚C)
2 NH 3 + 3Mg \u003d Mg 3 N 2 +3 H 2 (600 ˚C)
NH 3 (g) + CO 2 (g) + H 2 O \u003d NH 4 HCO 3 (room temperature, pressure)
Receipt. IN laboratories- displacement of ammonia from ammonium salts when heated with soda lime: Ca (OH) 2 + 2NH 4 Cl \u003d CaCl 2 + 2H 2 O + NH 3
Or boiling an aqueous solution of ammonia, followed by drying the gas.
In industry ammonia is produced from nitrogen with hydrogen. Produced by the industry either in liquefied form or in the form of a concentrated aqueous solution under the technical name ammonia water.

Ammonia hydrateNH 3 * H 2 O. Intermolecular connection. White, in crystal lattice are NH 3 and H 2 O molecules bound by a weak hydrogen bond. Present in aqueous solution ammonia, weak base(dissociation products - cation NH 4 and anion OH). The ammonium cation has a regular tetrahedral structure (sp 3 hybridization). Thermally unstable, completely decomposes when the solution is boiled. Neutralized by strong acids. Manifests restorative properties(due to N -3) in a concentrated solution. It enters into the reaction of ion exchange and complex formation.

Qualitative reaction– formation of white "smoke" upon contact with gaseous HCl. It is used to create a slightly alkaline environment in solution, during the precipitation of amphoteric hydroxides.
A 1 M ammonia solution contains mainly NH 3 *H 2 O hydrate and only 0.4% NH 4 OH ions (due to hydrate dissociation); thus, the ionic "ammonium hydroxide NH 4 OH" is practically not contained in the solution, there is no such compound in the solid hydrate either.
Equations of the most important reactions:
NH 3 H 2 O (conc.) = NH 3 + H 2 O (boiling with NaOH)
NH 3 H 2 O + HCl (diff.) = NH 4 Cl + H 2 O
3(NH 3 H 2 O) (conc.) + CrCl 3 = Cr(OH) 3 ↓ + 3 NH 4 Cl
8(NH 3 H 2 O) (conc.) + 3Br 2(p) = N 2 + 6 NH 4 Br + 8H 2 O (40-50˚C)
2(NH 3 H 2 O) (conc.) + 2KMnO 4 = N 2 + 2MnO 2 ↓ + 4H 2 O + 2KOH
4(NH 3 H 2 O) (conc.) + Ag 2 O = 2OH + 3H 2 O
4(NH 3 H 2 O) (conc.) + Cu(OH) 2 + (OH) 2 + 4H 2 O
6(NH 3 H 2 O) (conc.) + NiCl 2 = Cl 2 + 6H 2 O
A dilute ammonia solution (3-10%) is often called ammonia(the name was invented by alchemists), and a concentrated solution (18.5 - 25%) is an ammonia solution (produced by industry).

nitrogen oxides

nitrogen monoxideNO

Non-salt forming oxide. colorless gas. The radical contains a covalent σπ-bond (N꞊O), in the solid state the dimer N 2 O 2 co N-N connection. Extremely thermally stable. Sensitive to atmospheric oxygen (turns brown). Slightly soluble in water and does not react with it. Chemically passive in relation to acids and alkalis. When heated, it reacts with metals and non-metals. highly reactive mixture of NO and NO 2 ("nitrous gases"). Intermediate in synthesis nitric acid.
Equations of the most important reactions:
2NO + O 2 (ex.) = 2NO 2 (20˚C)
2NO + C (graphite) \u003d N 2 + CO 2 (400-500˚C)
10NO + 4P(red) = 5N 2 + 2P 2 O 5 (150-200˚C)
2NO + 4Cu \u003d N 2 + 2 Cu 2 O (500-600˚C)
Reactions to mixtures of NO and NO 2:
NO + NO 2 + H 2 O \u003d 2HNO 2 (p)
NO + NO 2 + 2KOH(razb.) \u003d 2KNO 2 + H 2 O
NO + NO 2 + Na 2 CO 3 \u003d 2Na 2 NO 2 + CO 2 (450-500˚C)
Receipt in industry: oxidation of ammonia with oxygen on a catalyst, in laboratories- interaction of dilute nitric acid with reducing agents:
8HNO 3 + 6Hg \u003d 3Hg 2 (NO 3) 2 + 2 NO+ 4 H 2 O
or reduction of nitrates:
2NaNO 2 + 2H 2 SO 4 + 2NaI \u003d 2 NO + I 2 ↓ + 2 H 2 O + 2Na 2 SO 4

nitrogen dioxideNO 2

Acid oxide, conditionally corresponds to two acids - HNO 2 and HNO 3 (acid for N 4 does not exist). Brown gas, monomer NO 2 at room temperature, liquid colorless dimer N 2 O 4 (dianitrogen tetroxide) in the cold. Completely reacts with water, alkalis. Very strong oxidizing agent, corrosive to metals. It is used for the synthesis of nitric acid and anhydrous nitrates, as a rocket fuel oxidizer, an oil cleaner from sulfur, and an oxidation catalyst. organic compounds. Poisonous.
The equation of the most important reactions:
2NO 2 ↔ 2NO + O 2
4NO 2 (l) + H 2 O \u003d 2HNO 3 + N 2 O 3 (syn.) (in the cold)
3 NO 2 + H 2 O \u003d 3HNO 3 + NO
2NO 2 + 2NaOH (diff.) \u003d NaNO 2 + NaNO 3 + H 2 O
4NO 2 + O 2 + 2 H 2 O \u003d 4 HNO 3
4NO 2 + O 2 + KOH \u003d KNO 3 + 2 H 2 O
2NO 2 + 7H 2 = 2NH 3 + 4 H 2 O (cat. Pt, Ni)
NO 2 + 2HI(p) = NO + I 2 ↓ + H 2 O
NO 2 + H 2 O + SO 2 = H 2 SO 4 + NO (50-60˚C)
NO 2 + K = KNO 2
6NO 2 + Bi(NO 3) 3 + 3NO (70- 110˚C)
Receipt: in industry - oxidation of NO with atmospheric oxygen, in laboratories– interaction of concentrated nitric acid with reducing agents:
6HNO 3 (conc., mountains) + S \u003d H 2 SO 4 + 6NO 2 + 2H 2 O
5HNO 3 (conc.,hort.) + P (red) \u003d H 3 PO 4 + 5NO 2 + H 2 O
2HNO 3 (conc., mountains) + SO 2 \u003d H 2 SO 4 + 2 NO 2

dinitrogen oxideN 2 O

Colorless gas with a pleasant smell ("laughing gas"), N꞊N꞊О, formal nitrogen oxidation state +1, poorly soluble in water. Supports the combustion of graphite and magnesium:

2N 2 O + C = CO 2 + 2N 2 (450˚C)
N 2 O + Mg = N 2 + MgO (500˚C)
receive thermal decomposition ammonium nitrate:
NH 4 NO 3 \u003d N 2 O + 2 H 2 O (195-245˚C)
used in medicine as an anesthetic.

dinitrogen trioxideN 2 O 3

At low temperatures, it is a blue liquid, ON꞊NO 2, the formal oxidation state of nitrogen is +3. At 20 ˚C, it decomposes by 90% into a mixture of colorless NO and brown NO 2 (“nitrous gases”, industrial smoke - “fox tail”). N 2 O 3 - acid oxide, forms HNO 2 with water in the cold, reacts differently when heated:
3N 2 O 3 + H 2 O \u003d 2HNO 3 + 4NO
With alkalis gives HNO 2 salts, for example NaNO 2 .
Obtained by the interaction of NO with O 2 (4NO + 3O 2 \u003d 2N 2 O 3) or with NO 2 (NO 2 + NO \u003d N 2 O 3)
with strong cooling. "Nitrous gases" and environmentally hazardous, act as catalysts for the destruction of the ozone layer of the atmosphere.

dinitrogen pentoxide N 2 O 5

Colorless, solid, O 2 N - O - NO 2, nitrogen oxidation state is +5. At room temperature, it decomposes into NO 2 and O 2 in 10 hours. Reacts with water and alkalis as an acidic oxide:
N 2 O 5 + H 2 O \u003d 2HNO 3
N 2 O 5 + 2NaOH \u003d 2NaNO 3 + H 2
Obtained by dehydration of fuming nitric acid:
2HNO 3 + P 2 O 5 \u003d N 2 O 5 + 2HPO 3
or oxidation of NO 2 with ozone at -78˚C:
2NO 2 + O 3 \u003d N 2 O 5 + O 2

Nitrites and nitrates

Potassium nitriteKNO 2 . White, hygroscopic. Melts without decomposition. Stable in dry air. Let's very well dissolve in water (forming colorless solution), it is hydrolyzed on anion. A typical oxidizing and reducing agent in an acidic environment, reacts very slowly in an alkaline environment. Enters into ion exchange reactions. Qualitative reactions on the NO 2 ion - discoloration of the violet solution of MnO 4 and the appearance of a black precipitate when I ions are added. It is used in the production of dyes, as an analytical reagent for amino acids and iodides, a component of photographic reagents.
equation of the most important reactions:
2KNO 2 (t) + 2HNO 3 (conc.) \u003d NO 2 + NO + H 2 O + 2KNO 3
2KNO 2 (dil.) + O 2 (ex.) → 2KNO 3 (60-80 ˚C)
KNO 2 + H 2 O + Br 2 = KNO 3 + 2HBr
5NO 2 - + 6H + + 2MnO 4 - (violet) \u003d 5NO 3 - + 2Mn 2+ (bts.) + 3H 2 O
3 NO 2 - + 8H + + CrO 7 2- \u003d 3NO 3 - + 2Cr 3+ + 4H 2 O
NO 2 - (saturated) + NH 4 + (saturated) \u003d N 2 + 2H 2 O
2NO 2 - + 4H + + 2I - (BC) = 2NO + I 2 (black) ↓ = 2H 2 O
NO 2 - (razb.) + Ag + \u003d AgNO 2 (light yellow) ↓
Receipt inindustry– recovery of potassium nitrate in the processes:
KNO 3 + Pb = KNO 2+ PbO (350-400˚C)
KNO 3 (conc.) + Pb (sponge) + H 2 O = KNO 2+ Pb(OH) 2 ↓
3 KNO 3 + CaO + SO 2 \u003d 2 KNO 2+ CaSO 4 (300 ˚C)

H itrat potassium KNO 3
technical name potassium, or indian salt , saltpeter. White, melts without decomposition, decomposes upon further heating. Air resistant. Highly soluble in water (high endo-effect, = -36 kJ), there is no hydrolysis. A strong oxidizing agent when fused (due to the release of atomic oxygen). In solution, it is reduced only by atomic hydrogen (in an acid medium to KNO 2, in an alkaline medium to NH 3). Used in glass production as a preservative food products, a component of pyrotechnic mixtures and mineral fertilizers.

2KNO 3 \u003d 2KNO 2 + O 2 (400-500 ˚C)

KNO 3 + 2H 0 (Zn, diluted HCl) = KNO 2 + H 2 O

KNO 3 + 8H 0 (Al, conc. KOH) = NH 3 + 2H 2 O + KOH (80 ˚C)

KNO 3 + NH 4 Cl \u003d N 2 O + 2H 2 O + KCl (230-300 ˚C)

2 KNO 3 + 3C (graphite) + S = N 2 + 3CO 2 + K 2 S (combustion)

KNO 3 + Pb = KNO 2 + PbO (350 - 400 ˚C)

KNO 3 + 2KOH + MnO 2 = K 2 MnO 4 + KNO 2 + H 2 O (350 - 400 ˚C)

Receipt: in industry
4KOH (horizontal) + 4NO 2 + O 2 = 4KNO 3 + 2H 2 O

and in the lab:
KCl + AgNO 3 \u003d KNO 3 + AgCl ↓