At ordinary temperatures, N 2 O is a colorless gas with a weak pleasant odor and a sweetish taste; has a narcotic effect, causing first convulsive laughter, then loss of consciousness.
Methods of obtaining
1. Decomposition of ammonium nitrate with slight heating:
NH 4 NO 3 = N 2 O + 2H 2 O
2. Effect of HNO 3 on active metals
10HNO 3 (conc.) + 4Ca = N 2 O + 4Ca(NO 3) 2 + 5H 2 O
Chemical properties
N 2 O exhibits neither acidic nor basic properties, i.e. it does not interact with bases, acids, or water (non-salt-forming oxide).
At T > 500 "C, it decomposes into simple substances. N 2 O is a very strong oxidizing agent. For example, it is capable of oxidizing sulfur dioxide in an aqueous solution to sulfuric acid:
N 2 O + SO 2 + H 2 O = N 2 + H 2 SO 4
NO - nitric oxide (II), nitrogen monoxide.
At ordinary temperatures, NO is a colorless, odorless gas, slightly soluble in water, very toxic (in high concentrations it changes the structure of hemoglobin).
Methods of obtaining
1. Direct synthesis from simple substances can only be carried out at very high T:
N 2 + O 2 = 2NO - Q
2. Production in industry (1st stage of HNO 3 production).
4NH 3 + 5O 2 = 4NO + 6H 2 O
3. Laboratory method - dilute action. HNO 3 for heavy metals:
8HNO 3 + 3Cu = 2NO + 3Cu(NO 3) 2 + 4H 2 O
Chemical properties
NO is a non-salt-forming oxide (like N2O). It has redox duality.
I. NO - oxidizing agent
2NO + SO 2 + H 2 O = N 2 O + H 2 SO 4
2NO + 2H 2 = N 2 + 2H 2 O (with explosion)
II. NO - reducing agent
2NO + O 2 = 2NO 2
10NO + 6KMnO 4 + 9H 2 SO 4 = 10HNO 3 + 3K 2 SO 4 + 6MnSO 4 + 4H 2 O
NO 2 - nitric oxide (IV), nitrogen dioxide
At normal temperatures, NO 2 is a red-brown poisonous gas with a pungent odor. It is a mixture of NO 2 and its dimer N 2 O 4 in a ratio of -1:4. Nitrogen dioxide is highly soluble in water.
Methods of obtaining
I. Industrial - NO oxidation: 2NO + O 2 = 2NO 2
II. Laboratory:
action of conc. HNO 3 for heavy metals: 4HNO 3 + Cu = 2NO 2 + Cu(NO 3) 2 + 2H 2 O
nitrate decomposition: 2Pb(NO 3) 2 = 4NO 2 + O 2 + 2PbO
Chemical properties
NO 2 - acid oxide, mixed anhydride of 2 acids
NO 2 reacts with water, basic oxides and alkalis. But the reactions do not proceed the same way as with ordinary oxides - they are always redox. This is explained by the fact that there is no acid with CO. (N) = +4, therefore NO 2, when dissolved in water, disproportionates with the formation of 2 acids - nitric and nitrous:
2NO 2 + H 2 O = HNO 3 + HNO 2
If dissolution occurs in the presence of O 2, then one acid is formed - nitric acid:
4NO 2 + 2H 2 O + O 2 = 4HNO 3
The interaction of NO 2 with alkalis occurs in a similar way:
in the absence of O 2: 2NO 2 + 2NaOH = NaNO 3 + NaNO 2 + H 2 O
in the presence of O 2: 4NO 2 + 4NaOH + O 2 = 4NaNO 3 + 2H 2 O
NO 2 is a very strong oxidizing agent
In terms of oxidizing ability, NO 2 is superior to nitric acid. C, S, P, metals and some organic substances burn in its atmosphere. In this case, NO 2 is reduced to free nitrogen:
10NO 2 + 8P = 5N 2 + 4P 2 O 5
2NO 2 + 8HI = N 2 + 4I 2 + 4H 2 O (a violet flame appears)
In the presence of Pt or Ni, nitrogen dioxide is reduced by hydrogen to ammonia:
2NO 2 + 7H 2 = 2NH 3 + 4H 2 O
NO 2 is used as an oxidizer in rocket fuels. When it interacts with hydrazine and its derivatives, a large amount of energy is released:
2NO 2 + 2N 2 H 4 = 3N 2 + 4H 2 O + Q
N 2 O 3 and N 2 O 5 are unstable substances
Both oxides have a pronounced acidic character and are anhydrides of nitrous and nitric acids, respectively.
N 2 O 3 as an individual substance exists only in the solid state below T pl. (-10 0 C).
With increasing temperature it decomposes: N 2 O 3 → NO + NO 2
N 2 O 5 at room temperature and especially in light decomposes so vigorously that it sometimes spontaneously explodes.
Due to the fact that nitrogen exhibits different valences in its compounds, this element is characterized by several oxides: dinitrogen oxide, mono-, tri-, di- and pentoxides of nitrogen. Let's look at each of them in more detail.
DEFINITION
Dianitrogen oxide(laughing gas, nitrous oxide) is a colorless gas that is thermally stable.
Poorly soluble in water. When strongly cooled, clarate N 2 O×5.75H 2 O crystallizes from the solution.
DEFINITION
Nitrogen monoxide It can exist either as a colorless gas or as a blue liquid.
In the solid state it is completely dimerized (N 2 O 2), in the liquid state - partially (≈ 25% N 2 O 2), in the gas - to a very small extent. Extremely thermally stable. Poorly soluble in water.
DEFINITION
Nitrogen trioxide is a thermally unstable blue liquid.
At room temperature, it decomposes 90% into NO and NO 2 and turns brown (NO 2), has no boiling point (NO evaporates first). In the solid state, it is a white or bluish substance with an ionic structure - nitrosyl nitrite (NO +)(NO 2 -). In gas it has a molecular structure of ON-NO 2.
DEFINITION
Nitrogen dioxide(foxtail) is a brown gas.
At temperatures above 135 o C it is a monomer, at room temperature it is a red-brown mixture of NO 2 and its dimer (dianitrogen tetroxide) N 2 O 4. In the liquid state the dimer is colorless, in the solid state it is white. It dissolves well in cold water (a saturated solution is bright green), completely reacting with it.
DEFINITION
Nitrogen pentoxide (nitric anhydride) is a white solid, colorless gas and liquid.
When heated, it sublimes and melts; at room temperature it decomposes in 10 hours. In the solid state it has an ionic structure (NO 2 +)(NO 3 -) - nitroyl nitrate.
Table 1. Physical properties of nitrogen oxides.
Obtaining Nitric Oxide
In laboratory conditions, dinitrogen oxide is obtained by carefully heating dry ammonium nitrate (1) or by heating a mixture of sulfamic and nitric (73%) acids (2):
NH 4 NO 3 = N 2 O + 2H 2 O (1);
NH 2 SO 2 OH + HNO 3 = N 2 O + H 2 SO 4 + H 2 O (2).
Nitrogen monoxide is produced by the interaction of simple substances nitrogen and oxygen at high temperatures (≈1300 o C):
N2 + O2 = 2NO.
In addition, nitric oxide (II) is one of the products of the reaction of dissolving copper in dilute nitric acid:
3Cu + 8HNO 3 = 3Cu(NO 3) 2 + 2NO + 4H 2 O.
When a mixture of gases consisting of nitrogen oxides (II) and (IV) is cooled to -36 o C, nitrogen trioxide is formed:
NO + NO 2 = N 2 O 3.
This compound can be obtained by the action of 50% nitric acid on arsenic (III) oxide (3) or starch (4):
2HNO 3 + As 2 O 3 = NO 2 + NO + 2HAsO 3 (3);
HNO 3 + (C 6 H 10 O 5) n = 6nNO + 6nNO 2 + 6nCO 2 + 11nH 2 O (4).
Thermal decomposition of lead(II) nitrate leads to the formation of nitrogen dioxide:
2Pb(NO3)2 = 2PbO + 4NO2 + O2.
The same compound is formed when copper is dissolved in concentrated nitric acid:
Cu + 4HNO 3 = Cu(NO 3) 2 + 2NO 2 + 2H 2 O.
Nitrogen pentoxide is obtained by passing dry chlorine over dry silver nitrate (5), as well as by the reaction between nitrogen oxide (IV) and ozone (6):
2Cl 2 + 4AgNO 3 = 2N 2 O 5 + 4AgCl + O 2 (5);
2NO 2 + O 3 = N 2 O 5 + O 2 (6).
Chemical properties of nitric oxide
Dianitrogen oxide is slightly reactive and does not react with dilute acids, alkalis, ammonia hydrate, or oxygen. When heated, it reacts with concentrated sulfuric acid, hydrogen, metals, and ammonia. Supports combustion of carbon and phosphorus. In ORR it can exhibit the properties of both a weak oxidizing agent and a weak reducing agent.
Nitrogen monoxide does not react with water, dilute acids, alkalis, or ammonia hydrate. Instantly adds oxygen. When heated, it reacts with halogens and other non-metals, strong oxidizing and reducing agents. Enters into complexation reactions.
Nitrogen trioxide exhibits acidic properties and reacts with water, alkalis, and ammonia hydrate. Reacts vigorously with oxygen and ozone, oxidizes metals.
Nitrogen dioxide reacts with water and alkalis. In OVR it exhibits the properties of a strong oxidizing agent. Causes corrosion of metals.
Nitrogen pentoxide exhibits acidic properties and reacts with water, alkalis, and ammonia hydrate. It is a very strong oxidizing agent.
Applications of Nitric Oxide
Dianitrogen oxide is used in the food industry (a propellant in the production of whipped cream), medicine (for inhalation anesthesia), and also as a main component of rocket fuel.
Nitrogen trioxide and dioxide are used in inorganic synthesis to produce nitric and sulfuric acids. Nitrogen (IV) oxide has also found use as a component of rocket fuel and mixed explosives.
Examples of problem solving
EXAMPLE 1
| Exercise | Nitric oxide contains 63.2% oxygen. What is the formula of the oxide. |
| Solution | The mass fraction of element X in a molecule of the composition NX is calculated using the following formula: ω (X) = n × Ar (X) / M (HX) × 100%. Let's calculate the mass fraction of nitrogen in the oxide: ω(N) = 100% - ω(O) = 100% - 63.2% = 36.8%. Let us denote the number of moles of elements included in the compound by “x” (nitrogen) and “y” (oxygen). Then, the molar ratio will look like this (the values of relative atomic masses taken from D.I. Mendeleev’s Periodic Table are rounded to whole numbers): x:y = ω(N)/Ar(N) : ω(O)/Ar(O); x:y= 36.8/14: 63.2/16; x:y= 2.6: 3.95 = 1: 2. This means that the formula for the compound of nitrogen and oxygen will be NO 2. This is nitric oxide (IV). |
| Answer | NO 2 |
EXAMPLE 2
| Exercise | Which gases are heavier and which are lighter than air and by how many times: carbon dioxide, nitrogen dioxide, carbon monoxide, chlorine, ammonia? |
| Solution | The ratio of the mass of a given gas to the mass of another gas taken in the same volume, at the same temperature and the same pressure is called the relative density of the first gas to the second. This value shows how many times the first gas is heavier or lighter than the second gas. The relative molecular weight of air is taken to be 29 (taking into account the content of nitrogen, oxygen and other gases in the air). It should be noted that the concept of “relative molecular mass of air” is used conditionally, since air is a mixture of gases. D air (CO 2) = M r (CO 2) / M r (air); D air (CO 2) = 44 / 29 = 1.52. M r (CO 2) = A r (C) + 2 × A r (O) = 12 + 2 × 16 = 12 + 32 = 44. D air (NO 2) = M r (NO 2) / M r (air); D air (NO 2) = 46 / 29 = 1.59. M r (NO 2) = A r (N) + 2 × A r (O) = 14 + 2 × 16 = 14 + 32 = 46. D air (CO) = M r (CO) / M r (air); D air (CO) = 28 / 29 = 0.97. M r (CO) = A r (C) + A r (O) = 12 + 16 = 28. D air (Cl 2) = M r (Cl 2) / M r (air); D air (Cl 2) = 71 / 29 = 2.45. M r (Cl 2) = 2 × A r (Cl) = 2 × 35.5 = 71. D air (NH 3) = M r (NH 3) / M r (air); D air (NH 3) = 17 / 29 = 0.57. M r (NH 3) = A r (N) + 3 ×A r (H) = 14 + 3 ×1 = 17. |
| Answer | Carbon dioxide, nitrogen dioxide and chlorine are 1.52 heavier than air, respectively; 1.59 and 2.45 times, and carbon monoxide and ammonia are 0.97 and 0.57 times lighter. |
|
Dina Kamilevna Gainullina— Candidate of Biological Sciences, researcher at the Department of Human and Animal Physiology, Faculty of Biology, Moscow State University. M. V. Lomonosova, specialist in the field of circulatory physiology. Area of scientific interests: features of regulation of the vascular system in early postnatal ontogenesis. |
Svetlana Ivanovna Sofronova— postgraduate student of the same department, studies the problems of hormonal regulation of endothelial nitric oxide synthesis. |
Olga Sergeevna Tarasova- Doctor of Biological Sciences, Professor of the same department and leading researcher at the Laboratory of Physiology of Muscular Activity of the State Scientific Center of the Russian Federation "Institute of Medical and Biological Problems of the Russian Academy of Sciences", a specialist in the field of blood circulation and the autonomic nervous system. Area of scientific interests is the interaction of systemic and local mechanisms of regulation of the cardiovascular system. |
The tone of blood vessels and the level of blood pressure in the body are regulated by the coordinated work of many systems and mechanisms, among which the vascular endothelium plays an important role. The secretion of nitric oxide (NO) is one of the key functions of endothelial cells, and doctors often associate their dysfunction in various diseases with a decrease in NO production. What are the modern ideas about the operation of this system? We will try to answer this question in our article.
Background
The layer of cells lining all blood and lymphatic vessels, as well as the cardiac cavities, was first described in 1847 by T. Schwann as a “distinct membrane,” which 18 years later was called endothelium by W. Gies. In relatively large vessels (arteries and veins), this layer serves as a barrier between blood and smooth muscle cells, and the walls of the smallest vessels, capillaries, are entirely built of endothelial cells. Their total number is very large: in the body of an adult, the total mass exceeds 1 kg!
In the 50-60s of the XX century. Scientists, armed with an electron microscope, described in detail the structure of the endothelium, but its role in regulating the functions of the cardiovascular system remained unclear. Until 1980, the endothelium was considered only a selectively permeable barrier between the blood and the vascular wall, although already at that time it was known that it was capable of secreting substances that prevent blood clotting.
The beginning of modern ideas about the functions of the endothelium was laid in 1980, when R. Farchgott and J. Zawadzki drew attention to its role in the regulation of vascular tone. In elegant experiments, the researchers showed that a substance such as acetylcholine causes relaxation of aortic preparations isolated from the body of a rabbit only in the presence of endothelium. This observation turned out to be so important that Farchgott later became one of the Nobel Prize laureates (1998). Nowadays, the endothelium-dependent vascular reaction in response to acetylcholine and other substances is described in a huge number of scientific works performed on a wide variety of arterial vessels - not only large, but also small ones that regulate the blood supply to organs (Fig. 1).
By 1986, it became clear that relaxation of vascular smooth muscle is caused by nitric oxide (NO), which is released from the endothelium under the influence of acetylcholine. How, in such a short time (only six years), was it possible to isolate NO from a long series of other candidates for the role of mediator between the endothelium and vascular smooth muscle? The fact is that 10 years before the famous work of Farchgott and Zawadzki, the vasodilating effect of NO was studied. Indeed, by that time, nitroglycerin (it serves as a source of NO molecules) had been treating angina pectoris resulting from spasms of the heart vessels for 100 years. The identity of the endothelial relaxing factor and NO was also established by such indicators as extreme instability (especially in the presence of reactive oxygen species), inactivation when interacting with hemoglobin and related proteins, as well as the ability to cause similar biochemical changes in vascular smooth muscle cells.
In the human and animal body, nitric oxide is one of the key endogenous regulators of the cardiovascular and other systems. In 1992, it was named molecule of the year, and the annual number of publications on its functions in the body today amounts to several thousand. The endothelium can be called a giant endocrine organ, in which the cells are not collected together, as in the endocrine glands, but are dispersed in vessels that penetrate all the organs and tissues of our body. Under normal physiological conditions, the endothelium is activated mainly mechanically: by shear stress created by blood flow, or by stretching of a vessel under blood pressure. In addition, endothelial cells can be activated by regulatory molecules, such as purine compounds (ATP and ADP), peptides (bradykinin, calcitonin gene-related peptide, substance P, etc.).
In addition to nitric oxide, endothelial cells synthesize other substances that affect vascular tone, tissue blood supply and blood pressure. Thus, NO assistants in relaxing blood vessels can be prostacyclin (prostaglandin I 2) and endothelial hyperpolarizing factor. The proportion of their participation depends on the sex and species of the animal, the type of vascular bed and the size of the vessel. For example, the effect of NO is stronger in relatively large vessels, and the hyperpolarizing factor - in smaller ones.
The endothelium produces not only vasodilator substances, but also vasoconstrictors: some prostaglandins, thromboxane, endothelin-1 and angiotensin II peptides, superoxide anion. In a healthy body, the secretory activity of the endothelium is aimed at the production of vasodilating factors. But in various diseases (systemic or pulmonary hypertension, myocardial ischemia, diabetes mellitus, etc.) or in a healthy body during aging, the secretory phenotype of the endothelium can change towards vasoconstrictor effects.
Despite the variety of regulatory mechanisms dependent on the endothelium, its normal function is most often associated with the ability to secrete NO. When the endothelium changes its properties during diseases, doctors call this condition endothelial dysfunction, implying a decrease in NO production. In connection with this importance of NO, we will consider modern ideas about its regulatory role, first in normal conditions, and then in some forms of vascular pathology.
Synthesis and regulation of NO in the endothelium
In nature, the synthesis of nitric oxide can occur through various pathways. Thus, in the troposphere it is formed from O 2 and N 2 under the influence of lightning discharges, in plants - due to the photochemical reaction between NO 2 and carotenoids, and in the body of animals - during the interaction of nitrites and nitrates with proteins containing metal atoms (for example, with hemoglobin ). All of these reactions occur without the participation of biological catalysts - enzyme proteins, so it is relatively difficult to control the speed. However, in the animal body, the main amount of NO as a regulator of physiological processes is formed under the action of special enzymes NO synthases (NOS), and the source of the nitrogen atom is the amino acid L-arginine [,].
There are several varieties (isoforms) of NO synthases, which are encoded by different genes. In 1990, the neuronal form of the enzyme (nNOS) was isolated from rat brain. A little later, inducible NOS (iNOS) was discovered in the cells of the immune system (macrophages), and endothelial NOS (eNOS) was discovered in the endothelium. Another isoform of NOS is localized in mitochondria; it regulates cellular respiration processes. Since a large number of cofactors are involved in NO synthesis, all enzyme isoforms have specific binding sites for them. Each NOS molecule consists of two identical halves. To combine them into a dimer, the cofactor tetrahydrobiopterin is required. With its deficiency, eNOS switches to the production of reactive oxygen species (superoxide anion and H 2 O 2), which can lead to damage to the endothelium and other cells of the vascular wall.
Two isoforms of the enzyme - eNOS and nNOS - are called constitutive because they are always present in cells and synthesize NO in relatively small quantities (compared to iNOS), and the activity of these isoforms is regulated by physiological stimuli. In contrast, iNOS is constantly synthesized only in some cells, for example, in macrophages, and in endothelial, nervous and many others it appears only in response to external, mainly inflammatory, stimuli (for example, elements of bacterial cell walls - bacterial lipopolysaccharides). Active iNOS produces NO 1000 times faster than eNOS and nNOS. Macrophages use these large amounts of NO to kill pathogens before destroying them.
Thus, the main NO synthase in the vascular wall is eNOS, and it is found mainly in the endothelium. Transcription of the eNOS gene in smooth muscle cells is prevented by special mechanisms, such as methylation of the “start” site. The synthase binds to the outer membrane of the endothelial cell in special invaginations, caveolae, where a large number of regulatory molecules (various ion channels and receptors) are concentrated. This “fixation” of the enzyme ensures its functional connection with receptors and channels, which facilitates the regulation of eNOS activity. The caveolin protein is localized in caveolae, which inhibits enzyme activity in the absence of stimulating stimuli.
The functional role of endothelial NO synthase depends on the number of molecules in the cell (the level of eNOS gene expression) and on its activity. It should be noted that the synthesis of new protein molecules is relatively long-term, so it is used to ensure long-term changes in NO production, for example, when adapting the vascular system to physical activity or to high-altitude hypoxia. To quickly control NO synthesis, other mechanisms are used, primarily changes in the intracellular concentration of Ca 2+, a universal regulator of cellular functions. Let us immediately note that such physiological regulation is characteristic only of eNOS and nNOS, while for iNOS (a Ca 2+ independent enzyme) it occurs mainly at the level of gene expression.
An increase in Ca 2+ concentration to a certain threshold level is an indispensable condition for the cleavage of endothelial NO synthase from caveolin and its transition to the active state. In addition to Ca 2+, phosphorylation, i.e., the covalent attachment of a phosphoric acid residue carried out by intracellular enzymes - protein kinases, is of great importance for the regulation of eNOS activity. Phosphorylation alters the ability of eNOS to be activated by calcium (Figure 2). Protein kinases attach phosphoric acid residues to strictly defined amino acid residues of the eNOS molecule, among which the most important are serine at position 1177 (Ser1177) and threonine at position 495 (Thr495). The Ser1177 site is considered the main site of eNOS activation. It is known that the degree of its phosphorylation increases rapidly under the influence of important regulatory factors: shear stress, bradykinin, vascular endothelial growth factor and estradiol. The main enzyme that carries out this process is Akt (another name is protein kinase B), but other kinases are also known that can activate eNOS (we will talk about them later).
Phosphorylation at the Thr495 site reduces enzyme activity. Such a negative effect can be enhanced under certain pathological conditions - oxidative stress, diabetes mellitus, etc. On the contrary, under some normal physiological influences, phosphate is removed (i.e., dephosphorylation of Thr495 occurs), due to which the affinity of eNOS for Ca 2+ increases and, consequently, its activity increases. Thus, the intensity of eNOS activity in endothelial cells can be dynamically regulated by Ca 2+ levels and phosphorylation/dephosphorylation by various protein kinases. This ultimately provides fine regulation of nitric oxide synthesis and, accordingly, its physiological effects on the cardiovascular system.
Mechanisms of relaxation of smooth muscle cells
How does NO secreted by endothelial cells cause vasodilation? The contraction of all types of muscle cells is ensured by the interaction of two proteins - actin and myosin, and the motor activity of the latter in smooth muscle cells appears only after its phosphorylation. This implies the presence of a large number of regulatory mechanisms that affect the contractile activity of smooth muscle cells, including nitric oxide.
NO molecules are lipophilic, so they freely penetrate from endothelial cells into smooth muscle cells. In them, the main NO acceptor is the enzyme guanylate cyclase, located in the cytosol and therefore called soluble (i.e., not associated with cell membranes). Guanylate cyclase, activated by nitric oxide, synthesizes cyclic guanosine monophosphate (cGMP), which serves as a potent activator of another enzyme, protein kinase G. Its targets in smooth muscle cells are numerous proteins involved in the regulation of cytoplasmic Ca 2+ concentration.
Protein kinase G activates certain types of potassium channels, which causes hyperpolarization (a shift in membrane potential towards more negative values) of smooth muscle cells, closes voltage-gated calcium channels of the outer membrane and thereby reduces the entry of Ca 2+ into the cell. In addition, this enzyme, in its active state, suppresses the release of Ca 2+ from intracellular stores and also promotes its removal from the cytoplasm. This also reduces the concentration of Ca 2+ and relaxes smooth muscles.
In addition to influencing Ca 2+ homeostasis, protein kinase G regulates Ca 2+ sensitivity of the contractile apparatus of smooth muscle cells, i.e., it reduces its ability to be activated when Ca 2+ increases. It is known that activation of protein kinase G (with the participation of intermediaries) reduces the level of phosphorylation of smooth muscle myosin, as a result of which it interacts less well with actin, which promotes relaxation. The combination of the described events leads to vasodilation, increased blood flow in organs and a decrease in blood pressure.
Physiological regulation of NO production
The ability to produce NO serves as a marker of the normal functional state of the endothelium: eliminating the effects of NO in a healthy body (for example, by pharmacological blockade of eNOS) leads to vasoconstriction and an increase in systemic blood pressure. As a result of the action of almost all normal physiological stimuli, the content of NO synthase in the endothelium (and/or its activity) increases. The key factor regulating NO production is blood flow. As it moves through the vessel, shear stress occurs on the surface of the endothelium. This stimulus is transmitted to intracellular endothelial NO synthase through activation of mechanosensitive channels and Ca 2+ entry. Another transmission option is through membrane enzymes, if the activity of the protein kinase Akt increases and eNOS is phosphorylated (at the Ser1177 site). The blood flow ensures constant secretion of small amounts of NO by the endothelium (Fig. 3).
The glycocalyx plays an important role in the sensitivity of the endothelium to shear stress. This is a layer of polymer molecules of a carbohydrate nature covering cells, the thickness of which can be several micrometers and even exceed the thickness of the endothelium itself. Since the “bushes” of glycoproteins grow inside the lumen of the vessel, they are the ones who first experience the effect of blood flow. When deformed, the glycocalyx fibers transmit a signal to membrane proteins and then to eNOS. Although this mechanism has so far been little studied, its importance is evidenced by the fact that impaired vascular response to shear stress in various diseases (atherosclerosis, diabetes mellitus, etc.) is associated with “balding” of the endothelium, i.e., with a decrease in thickness and a change in structure glycocalyx.
An increase in blood flow velocity leads to the activation of endothelial NO synthase and to vasodilation, and such prolonged or repeated exposures increase the content of this enzyme in the endothelium. The beneficial effects of physical exercise are based on this: it is known that with the help of training you can significantly improve the functioning of the endothelium without the use of drugs! However, it should be noted that not all exercises have such a beneficial effect. Firstly, the load must be accompanied by an increase in the speed of blood flow in the working muscles, as happens with fast walking, running or cycling, and strength exercises with weights do not have such an effect. Secondly, you should not train through force: with excessive loads, the secretion of the main stress hormone, cortisol, sharply increases, which reduces the activity of eNOS.
Additional activation of endothelial NO synthase during physical exercise is provided by protein kinase activated by adenosine monophosphate (AMP), which is found in almost all cells of our body, including endothelial cells. This enzyme is called a “cell energy status sensor” because it is activated when the AMP/ATP ratio in the cell cytoplasm increases, i.e., energy consumption begins to exceed its production. In the endothelium of arteries located inside intensely contracting skeletal muscles, this can occur as a result of hypoxia - muscle cells consume a lot of O 2, and the vascular endothelium lacks it. In addition, it has recently been shown that activation of this protein kinase in endothelial cells is possible with an increase in shear stress, i.e., with increased blood flow to working muscles. Activated protein kinase phosphorylates eNOS at the Ser1177 site, NO production increases and blood vessels dilate.
Cardiologists are well aware that through regular physical training it is possible to improve endothelial function not only in skeletal muscles and the heart, which are intensively supplied with blood during work, but also in organs not directly involved in training - in the brain, skin, etc. d. This suggests that in addition to the influence of blood flow on the endothelium, there are other mechanisms for the regulation of endothelial NO synthase. Among them, the leading role belongs to hormones, which are produced by the endocrine glands, transported by the blood and recognize target cells in various organs by the presence of special receptor proteins.
Among the hormones that can affect endothelial function during physical activity, we note growth hormone (somatotropic hormone), which is secreted by the pituitary gland. Both by itself and through its intermediaries, insulin-like growth factors, growth hormone increases the formation of endothelial NO synthase and its activity.
The most famous example of hormonal regulation of endothelial functions is the influence of female sex hormones, estrogens. Initially, this idea was formed thanks to epidemiological observations, when it turned out that for some reason women of childbearing age, compared to men, suffer less from vascular disorders associated with endothelial dysfunction. Moreover, in women, its ability to produce NO changes during the menstrual cycle, and in the first half, when the concentration of estrogen in the blood is high, endothelium-dependent vasodilation is more pronounced. These observations prompted numerous animal experiments. Thus, removal of the ovaries from female rats reduced the content and activity of endothelial NO synthase in the arteries of various organs (brain, heart, skeletal muscles, kidneys, intestines, etc.), and the administration of estrogens to such females contributed to the normalization of the impaired function. The effect of estrogens on eNOS activity is associated with activation of the protein kinase Akt, and the increase in eNOS synthesis is associated with their effect on the genome of endothelial cells.
It is interesting that disturbances in the reactions of the brain arteries were also found in experiments with the removal of the gonads in males, although the testes do not secrete estrogens, but androgens, male sex hormones. This paradox became clear when aromatase, an enzyme that converts androgens into estrogens, was discovered in the endothelium of brain arteries. Thus, the protective effect of estrogens on the vascular endothelium may also occur in males. However, in this case we should talk about local regulation, which is provided by estrogens formed directly in the vascular wall.
In conclusion, we will consider the regulation of endothelial NO synthase by thyroid hormones. It is known that when its functioning is disrupted in the vascular endothelium, the intensity of NO synthesis changes: in hyperthyroidism it increases, and in hypothyroidism it decreases. This effect is mainly due to changes in the content of NO synthase in endothelial cells. However, recently there has been evidence of the existence of another mechanism of action of these hormones on vascular endothelial cells. Thus, the Ca 2+ -dependent activity of eNOS and the degree of its phosphorylation at the Ser1177 site in the arteries of rats with experimental hyperthyroidism turned out to be significantly higher than in rats with hypothyroidism.
Thyroid hormones are known to play a key role in tissue differentiation in the developing organism. But their influence is not limited to just accelerating or slowing down processes, but often has a programming character. This means that if there is a lack of thyroid hormones at a certain critical age, the cells will not be able to become fully functioning, even if hormones are administered at later stages of life (in humans, hormonal therapy is effective only during the first months after birth). The mechanisms of the programming influence of thyroid hormones have been studied in detail only for the nervous system, and for other systems - much less well. However, it is well known that maternal hypothyroidism during pregnancy is, among other things, a risk factor for the development of cardiovascular diseases in the child. Interestingly, in the arteries of rat pups in the first weeks after birth, increased levels of thyroid hormone receptors are detected, as well as the enzyme deiodinase, which converts thyroxine (tetraiodothyronine) into the more active triiodothyronine. Based on these observations, it is tempting to assume that thyroid hormones may also have a programming effect on the vascular endothelium. Future research will show how true this is.
Mechanisms of impaired NO secretion by the endothelium
Unfortunately, the ability of our vascular endothelium to produce NO is not unlimited. The activity of the body's regulatory systems is high at a young and mature age, but decreases with aging under the influence of a number of factors. Firstly, few older people can emulate the saying of the ancient Greek philosopher Aristotle: “Life requires movement.” Secondly, with age, the activity of many hormonal systems fades: the secretion of growth hormone and sex hormones decreases, and the thyroid gland “falls asleep.” Thirdly, changes occur in the metabolism of all cells. In particular, the energy stations of the cell, mitochondria, begin to produce large quantities of reactive oxygen species, which inactivate NO, and also suppress the activity and reduce the content of endothelial NO synthase. It appears that age-related changes in the endothelium cannot be prevented, but they can be slowed by increasing mobility, limiting the intake of high-calorie foods (this also increases the activity of AMP-activated protein kinase), using hormone replacement therapy (for example, in postmenopausal women) or antioxidants, which have been developed and remains a priority area of gerontology.
Why is NO synthesis in the vascular endothelium disrupted in various pathologies? Two types of changes are possible here: rapid (decreased NO synthase activity in the endothelium), and long-term - a decrease in its content in cells. We will not consider various diseases separately, but will list their common mechanisms of harmful influence on the functioning of eNOS. A decrease in the activity of this enzyme in diseases is usually associated with an increase in its phosphorylation at the Thr495 site, caused by an increase in the activity of protein kinase C. Its powerful activator is diacylglycerol. Normally, it is a secondary messenger in signal transmission from many membrane receptors, but its excessive accumulation in endothelial cells leads to pathology.
A striking example of such changes can be a disease such as diabetes, in which a violation of the synthesis or action of insulin on cells leads to an increased concentration of glucose in the blood. Since glucose transport into the endothelium is not regulated by insulin (unlike cells of skeletal muscles, heart, adipose tissue and some others), sugar accumulates there and becomes a substrate for the synthesis of diacylglycerol, which activates protein kinase C.
The already mentioned oxidative stress is a marker of many cardiovascular pathologies. Increased formation of reactive oxygen species is characteristic of diabetes mellitus, atherosclerosis, and many forms of arterial hypertension. In these conditions, high activity of the renin-angiotensin system is often observed, and angiotensin II is a powerful provocateur of oxidative stress, which, on the one hand, reduces eNOS activity (for example, oxidized low-density lipoproteins can activate protein kinase C), and on the other, reduces gene expression eNOS, which also reduces NO production. The use of antioxidants or substances that interfere with the formation or action of angiotensin II (angiotensin-converting enzyme inhibitors or angiotensin II blockers) almost always increases the formation of NO. It must be said that a decrease in nitric oxide production in diseases may be associated not only with a direct effect on eNOS. Thus, the effect of glucocorticoids on the endothelium reduces the content of not only the enzyme itself, but also its cofactor, tetrahydrobiopterin.
Impaired functioning of endothelial NO synthase may be due to a lack of its main substrate, L-arginine. As a rule, this amino acid enters the body with food in sufficient quantities and, in addition, can be directly synthesized in the adult body. However, in addition to NO synthases, arginine serves as a substrate for many other enzymes, in particular arginase, which is located in various types of cells, including the vascular endothelium. In diabetes mellitus, oxidative stress, as well as in inflammatory processes under the influence of cytokines secreted by cells of the immune system (tumor necrosis factor, etc.), the content of arginase in the endothelium increases.
Finally, endothelial NO synthase inhibitors, such as dimethylarginine, may appear in humans and other animals. This “false substrate” of endothelial NO synthase competes with the true substrate, L-arginine, for the active site of the enzyme. Normally, dimethylarginine is formed in the body only in small quantities (in an adult ~60 mg/day), however, with a variety of circulatory pathologies (arterial hypertension, atherosclerosis, coronary insufficiency, etc.), its production increases significantly, and the activity of endothelial NO -synthase, accordingly, decreases.
So, nitric oxide is an important regulatory factor through which the endothelium has a relaxing effect on neighboring smooth muscle cells, causing vasodilation and smoothing out unwanted increases in blood pressure at the systemic level. As long as the endothelium retains the ability to secrete NO in quantities sufficient to solve these problems, there is no need to worry about the state of the vascular system.
This work was supported by the Russian Foundation for Basic Research. Project NK 14-04-31377 mol-a.
Literature
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Furchgott R. F., Zawadzki J. V. The obligatory role of endothelial cells in the relaxation of arterial smooth muscle by acetylcholine // Nature. 1980. V. 288. P. 373–376.
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Melkumyants A. M., Balashov S. A. Mechanical sensitivity of the arterial endothelium. Tver, 2005.
Calcitonin Gene Related Peptide) is formed from the same gene as calcitonin by alternative mRNA splicing in the brain and peripheral nervous system.
Residue numbers are given according to location in the human eNOS molecule.
Nitrogen forms a series of oxides with oxygen; all of them can be obtained from nitric acid or its salts.
Nitric oxide(I), or nitrous oxide, N 2 O is obtained by heating ammonium nitrate:
Nitric oxide (1) is a colorless gas with a faint odor and a sweetish taste. It is slightly soluble in water: one volume of water at 20 °C dissolves 0.63 volumes of N 2 O.
Nitric oxide (I) is a thermodynamically unstable compound. The standard Gibbs energy of its formation is positive (DS°b p =
104 kJ/mol). However, due to the high strength of bonds in the N 2 O molecule, the activation energies of reactions occurring with the participation of this substance are high. In particular, the activation energy for the decomposition of N 2 O is high. Therefore, at room temperature, nitrogen oxide (I) is stable. However, at elevated temperatures it decomposes into nitrogen and oxygen; Decomposition occurs the faster the higher the temperature.
Nitric oxide (1) does not react with water, acids or alkalis.
The electronic structure of the N 2 O molecule is discussed in § 41.
Inhalation of small amounts of nitric oxide (I) leads to dulling of pain sensitivity, as a result of which this gas is sometimes used in a mixture with oxygen for anesthesia. Large amounts of nitric oxide (I) have a stimulating effect on the nervous system; That's why it used to be called "laughing gas."
Nitric oxide(II), or nitric oxide, NO is a colorless, difficult to liquefy gas. Liquid nitric oxide (II) boils at -151.7°C and solidifies at -163.7°C. It is slightly soluble in water: 1 volume of water dissolves only 0.07 volumes of NO at 0°C.
According to its chemical properties, nitric oxide (II) is one of the indifferent oxides, since it does not form any acid.
Like N 2 O, nitrogen (II) oxide is thermodynamically unstable - the standard Gibbs energy of its formation is positive (AGo 6p = 86.6 kJ/mol). But, again, like N 2 O, NO does not decompose at room temperature because its molecules are quite strong. Only at temperatures above 1000 0 C does its decomposition into nitrogen and oxygen begin to occur at a noticeable rate. At very high temperatures, for reasons discussed in § 65, the decomposition of NO does not proceed completely - equilibrium is established in the NO-N 2 -O 2 system. Thanks to this, nitrogen oxide (II) can be obtained from simple substances at electric arc temperatures (3000-4000 ° C).
In the laboratory, nitric oxide (II) is usually obtained by reacting 30-35% nitric acid with copper:
In industry, it is an intermediate product in the production of nitric acid (see § 143).
Nitric oxide (II) is characterized by redox duality. Under the influence of strong oxidizing agents, it is oxidized, and in the presence of strong reducing agents, it is reduced. For example, it is easily oxidized by atmospheric oxygen to nitrogen dioxide: 
At the same time, a mixture of equal volumes of NO and H2 explodes when heated:
The electronic structure of the NO molecule is best described by the MO method. In Fig. Figure 116 shows a diagram of filling MO in a NO molecule (with similar diagrams for N 2 and CO molecules - see Fig. 51 and 53). The NO molecule has one more electron than the N 2 and CO molecules: this electron is located in the antibonding orbital l dist 2 R. Thus, the number of bonding electrons here exceeds the number of antibonding electrons by five. This corresponds to a coupling factor of 2.5 (5:2 = 2.5). Indeed, the dissociation energy of the NO molecule into atoms (632 kJ/mol) has an intermediate value compared to the corresponding values for the O 2 molecule (498 kJ/mol), in which the bond multiplicity is two, and the N 2 molecule (945 kJ/mol) , where the bond is triple. At the same time, in terms of dissociation energy, the NO molecule is close to the molecular oxygen ion O 2 (644 kJ/mol), in which the bond multiplicity is also 2.5.
When one electron is removed from a NO molecule, an NO + ion is formed, which does not contain antibonding electrons; the multiplicity of bonds between atoms increases to three (six bonding electrons). Therefore, the dissociation energy of the NO + ion (1050 kJ/mol) is higher than the dissociation energy of the NO molecule and is close to the corresponding value for the CO molecule (1076 kJ/mol), in which the bond multiplicity is three.
Rice. 116.
Dioxide(or nitrogen dioxide) NO 2 is a brown poisonous gas with a characteristic odor. It easily condenses into a reddish liquid (boiling point 21 0 C), which upon cooling gradually becomes lighter and freezes at -11.2 ° C, forming a colorless crystalline mass. When gaseous nitrogen dioxide is heated, its color, on the contrary, intensifies, and at 140 °C it becomes almost black. A change in the color of nitrogen dioxide with increasing temperature is accompanied by a change in its molecular weight. At low temperatures, the vapor density approximately corresponds to twice the formula N 2 O 4. With increasing temperature, the vapor density decreases and at 140 °C it corresponds to the formula NO 2. Colorless crystals, existing at -11.2 0 C and below, consist of N 2 O 4 molecules. As the N 2 O 4 molecules heat up, they dissociate to form dark brown nitrogen dioxide molecules; complete dissociation occurs at 140 0 C. Thus, at temperatures from -11.2 to 140 ° C, NO 2 and N 2 O 4 molecules are in equilibrium with each other:
Above 140 °C, the dissociation of NO 2 into NO and oxygen begins.
Nitrogen dioxide is a very energetic oxidizing agent. Many substances can burn in an NO 2 atmosphere, removing oxygen from it. It oxidizes sulfur dioxide into trioxide, which is what the nitrous method of producing sulfuric acid is based on (see § 131).
NO 2 vapors are poisonous. Inhaling them causes severe irritation of the respiratory tract and can lead to serious poisoning.
When dissolved in water, NO 2 reacts with water, forming nitric and nitrous acids:
But nitrous acid is very unstable and quickly decomposes:
Therefore, in practice, the interaction of nitrogen dioxide with water, especially hot water, proceeds according to the equation
which can be obtained by adding the two previous equations, if you first multiply the first of them by three.
In the presence of air, the resulting nitric oxide is immediately oxidized to nitrogen dioxide, so that in this case the NO 2 ends up completely converted to nitric acid:
This reaction is used in modern methods for producing nitric acid.
If nitrogen dioxide is dissolved in alkalis, a mixture of salts of nitric and nitrous acids is formed, for example:
Nitric oxide(III), or nitrous anhydride, N 2 O 3 is a dark blue liquid that decomposes into NO and NO 2 even at low temperatures. A mixture of equal volumes of NO and NO 2 upon cooling again forms N 2 O 3:
Nitrous acid HNO 2 corresponds to nitrogen oxide (III).
Nitric oxide(V), or nitric anhydride, N 2 O 5 are white crystals that gradually decompose into NO 2 and O 2 even at room temperature. It can be prepared by the action of phosphoric anhydride on nitric acid:
Nitric oxide (V) is a very strong oxidizing agent. Many organic substances ignite upon contact with it. In water, nitric oxide (V) dissolves well to form nitric acid.
In the solid state, N 2 O 5 is formed by the nitrate ion NO 3 and the ion
nitronium NO2. The latter contains the same number of electrons as can
CO 2 molecule and, like the latter, has a linear structure: O=N=O.
In pairs, the N 2 O 5 molecule is symmetrical; its structure can be represented by the following valence scheme, in which three-center bonds are shown with a dotted line (compare with the valence scheme of the nitric acid molecule).
Nitrogen oxides and hydroxidesTHE SECRET OF OXIDATION DEGREES
Nitrogen forms a number of oxides that formally correspond to all possible oxidation states from +1 to +5: N 2 O, NO, N 2 O 3, NO 2, N 2 O 5, but only two of them are nitric oxide (II) and oxide nitrogen(IV) - not only stable under normal conditions, but also actively involved in natural and industrial nitrogen cycles. Therefore, we will study their properties (in comparison). Let's start, as usual, with the structure of molecules.
Structure of nitrogen oxide molecules
Molecule NO. The structure is quite simple to assume: oxygen has two unpaired electrons, nitrogen has three - a double bond is formed and one unpaired electron in the remainder... It is not easy to answer the question why such a “non-standard” molecule is stable. By the way, it is worth noting that stable free radicals - molecules with unpaired electrons - are quite rare in nature. It can be assumed that NO molecules will pair and form a doubled, or dimeric, ONNO molecule. This way we can solve the problem of the unpaired electron.
Molecule NO2. It would seem that it couldn’t be simpler - an oxygen atom was attached to the NO molecule via an unpaired electron. (In fact, it is not an atom that is attached, but a molecule, and not to NO, but to the ONNO dimer. That is why the rate of addition decreases with increasing temperature - the dimer breaks into halves.) And now oxygen has an unpaired electron - a nitric oxide molecule (IV) is also a free radical. However, it is known that when two NO 2 molecules join and form an N 2 O 4 molecule, the connection occurs through nitrogen atoms, which means that it is nitrogen that should have this very unpaired electron. How can this be done?
The answer is unconventional, but quite in the “character” of nitrogen – a donor-acceptor bond. Using logic, consider the electrons that the nitrogen atom in the NO molecule has. This is an unpaired electron, a free pair of electrons and two more electrons bonded to oxygen - a total of five. And the oxygen atom “coming into contact” has six electrons in four orbitals. If you arrange them in twos, then one orbital will remain free. It is precisely this that is occupied by a pair of electrons of the nitrogen atom, and the unpaired electron in this connection turns out to have absolutely nothing to do with it (Fig. 1, 2).
It is worth mentioning one more point - since a pair of electrons located on s-orbital, “got into contact”, it was simply obliged to undergo hybridization - it is very difficult to offer the second atom for common use a pair of electrons, evenly distributed over the surface of the first atom. The question arises: what type of hybridization does the atom use? Answer: the three electron orbitals of nitrogen are in the state sp 2-hybridization. The NO 2 molecule is angular, the angle is 134° (the angle is greater than 120° because one electron repels bond electrons from itself weaker than a pair of electrons) (Fig. 3–5).
Physical properties of nitrogen oxides
Nitric oxide(II) NO. Molecular crystal lattice; the molecule is light, weakly polar (the electronegativity of oxygen is higher than that of nitrogen, but not by much). It can be assumed that the melting and boiling points will be low, but higher than those of nitrogen, since any polarity of the molecule makes it possible to connect electrostatic forces of attraction to simply intermolecular forces. The formation of a dimer also increases the boiling point, making the molecule heavier. The structure of the molecule also suggests low solubility in water, a solvent noticeably more polar than NO. It is worth emphasizing that nitrogen(II) oxide is neither color nor odor.
Nitric oxide(IV) NO2. The crystal lattice is also molecular, but since the molecule itself is heavier than NO and its tendency to dimerize is noticeably higher, this substance should melt and boil at noticeably higher temperatures. The boiling point is 21 °C, so under normal conditions it is 20 °C and 760 mm Hg. Art. – nitrogen(IV) oxide, liquid.
Now let's look at solubility. Let us remember that the word “solubility” can also mean chemical reactions with water; the main thing is that the solvent absorbs what is being dissolved. When oxides react with water, as is known, hydroxides are obtained - formally these are just hydrated oxides, but reality often presents a lot of interesting and completely informal things. So this nitric oxide dissolves in water, reacting with it at the same time, and thus two acids are obtained at once!
Note that nitric oxide (IV) has a characteristic pungent odor and a reddish-brown color, the shades of which differ from each other depending on the concentration. It is for this color that emissions of nitrogen oxides into the atmosphere are called “fox tails.”
You ask: where is the secret? The first part of the mystery of oxidation states is in front of you: why are oxides with even oxidation states stable for an element of the fifth (odd) group? (At the same time, there are also free radicals!) In the most general sense, the answer is obvious - since they are stable, it means that it is so beneficial for them. Energetically. And why? Apparently, the point is in the specific structure of nitrogen and oxygen atoms - they have too many electrons and too few orbitals. It is “orbital capabilities” that dictate their rules and establish such “energy benefits”. Then the numbers “two” and “four” become clear: two electrons are missing from oxygen to eight and both atoms have only four orbitals.
You can also say that NO is just... waiting for an oxygen molecule to turn into NO 2. Using a metaphor, we note that the “meaning of life” of many atoms is the desire to find a “life partner” - an atom or atoms of another element. Although there are, of course, “convinced bachelors” like gold.
Chemical properties of nitrogen oxides
1. Reactions with metals. Since the nitrogen atom in positive oxidation states is an oxidizing agent, and the higher the oxidation state, the stronger the ability to take electrons from other atoms, then nitrogen oxides will react with metals - essentially reducing agents. The resulting products can be completely different, depending on the reaction conditions and the metal itself. For example, to hot copper, all nitrogen oxides give up oxygen, and themselves turn into the simple substance nitrogen:
By the amount of copper oxide and nitrogen oxide formed, it is possible to determine which nitrogen oxide reacted with copper.
2. Reactions with non-metals. First of all, let's look at reactions with oxygen. Here there is a difference between the oxides, and a very significant one.
Oxide NO reacts with oxygen to form nitric oxide (IV). The reaction is reversible. Moreover, with increasing temperature, the rate of this reaction decreases:
2NO + O 2 = 2NO 2.
Oxide NO 2 does not react with oxygen at all.
Ozone converts both oxides into nitrogen oxide (V).
Nitric oxide(II) NO adds ozone completely:
2NO + O 3 = N 2 O 5.
Nitrogen oxide (IV) NO 2 in reaction with ozone also releases oxygen:
2NO 2 + O 3 = N 2 O 5 + O 2.
3. Reactions with water. NO oxide does not react with water. NO 2 oxide with water forms two acids - nitric (nitrogen oxidation state +5) and nitrous (nitrogen oxidation state +3). In the presence of oxygen, NO 2 oxide is completely converted into nitric acid:
2NO 2 + H 2 O = HNO 3 + HNO 2,
4NO 2 + O 2 + 2H 2 O = 4HNO 3.
4. Reactions with acids. None of the oxides - NO or NO 2 - reacts with acids.
5. Reactions with alkalis. Both nitrogen oxides react with alkalis.
Oxide NO forms a salt of nitrous acid, nitric oxide (I) and nitrogen with alkali:
10NO + 6NaOH = 6NaNO2 + N2O + N2 + 3H2O.
Oxide NO 2 forms salts of two acids with alkali - nitric and nitrous:
2NO 2 + 2NaOH = NaNO 3 + NaNO 2 + H 2 O.
Let's return to our mystery of oxidation states. During the transition of oxygen compounds of nitrogen from the “gas” state, where you can move freely, to the “aqueous solution” state, where there is more hustle and bustle, where collectivism flourishes, where polar water molecules exist and actively act, no one will allow the molecule, atom or ion to be alone, a “change of orientation” occurs. It is odd oxidation states that become stable, as befits an element from an odd group. (Stable, however, is relatively. Nitrous acid, for example, can only exist in solution, otherwise it decomposes. But acids formally corresponding to oxides of nitrogen (II) and (IV) do not exist at all. Everything is known by comparison.)
It is interesting that not only the clearly acidic oxide NO 2 reacts with alkalis, but also NO - non-acidic in properties and oxidation state, and compounds of other oxidation states - odd ones - are obtained! Secret? Quite!
The structure of the molecule of nitrogen(V) hydroxide - nitric acid
Of the nitrogen hydroxides, we will consider one, but the most large-tonnage one - nitric acid.
The nitric acid molecule is polar (primarily due to the different electronegativity of oxygen and hydrogen, because nitrogen is hidden inside the molecule) and asymmetric. All three angles between nitrogen and oxygen bonds present in it are different. The formal oxidation state of nitrogen is the highest, i.e. +5. But at the same time, the nitrogen atom has only four bonds with other atoms - the valency of nitrogen is four. Another secret.
It is clear how it could happen that valence of an atom numerically greater than its oxidation state. To do this, it is enough to form a bond between identical atoms in the molecule. For example, in hydrogen peroxide, oxygen has a valency of two, but the oxidation state is only –1. Oxygen managed to pull the common electron bond pair with hydrogen closer to itself, and the bond pair of two oxygen atoms remains strictly in the middle. But how to make it so that valence of an atom was less oxidation state?
Let's think: how does a nitric acid molecule actually work? The structure of a molecule is easier to understand if we consider the process of its preparation. Nitric acid is obtained by the reaction of nitrogen oxide (IV) with water (in the presence of oxygen): two NO 2 molecules simultaneously “attack” a water molecule with their unpaired electrons, as a result, the bond between hydrogen and oxygen is not broken as usual (a pair of electrons from oxygen and “naked proton"), and “honestly” - one NO 2 molecule gets hydrogen with its electron, the other gets the OH radical (Fig. 6). Two acids are formed: both acids are strong, both quickly give up their proton to the nearest water molecules and ultimately remain in the form of ions and . The ion is unstable, two HNO 2 molecules decompose into water, NO 2 and NO. NO oxide reacts with oxygen, turning into NO 2, and so on until only nitric acid is obtained.
Formally, it turns out that the nitrogen atom is connected to one oxygen atom by a double bond, and to the other by an ordinary single bond (this oxygen atom is also connected to a hydrogen atom). Nitrogen in HNO 3 is connected to the third oxygen atom by a donor-acceptor bond, with the nitrogen atom acting as a donor. The hybridization of the nitrogen atom should be sp 2 due to the presence of a double bond, which determines the structure - a flat triangle. In reality, it turns out that indeed a fragment of a nitrogen atom and three oxygen atoms is a flat triangle, only in a nitric acid molecule this triangle is incorrect - all three ONO angles are different, therefore, the sides of the triangle are different. When the molecule dissociates, the triangle becomes regular, equilateral. This means that the oxygen atoms in it become equivalent! All bonds become identical (a double bond is shorter than a single bond). How?
Let's reason. sp 2-Hybridization of the nitrogen atom forces the oxygen atoms to the same type of hybridization. The result is a flat structure, across which the p-orbitals that are not involved in hybridization, which are present in all four atoms, are located.
Now let's look at the total number of valence electrons: the ion contains five electrons from nitrogen, six from each of the three oxygen atoms, and another one that imparts a charge to the ion as a whole, for a total of twenty-four. Of these, six electrons are required to form three single bonds, twelve electrons are located along the perimeter of the molecule in hybrid orbitals (two electron pairs for each oxygen atom), leaving six electrons for the same four R-orbitals not involved in hybridization. The only reasonable explanation possible in this case is that all atoms share their electrons into a single electron cloud (Fig. 7). This is facilitated by small atomic radii and small interatomic distances. And symmetry is usually energetically favorable and therefore increases the stability of the structure as a whole. This is not the only case of the sharing of electrons by several atoms; similar “collective electron farming” is found in organic chemistry, for example, in aromatic compounds.
Let us return, however, to predictions of the properties of nitric acid based on ideas about the structure of the molecule. The obvious advantage of being in the form of an ion explains not only the high degree of dissociation of the acid in an aqueous solution, but also the possibility of dissociation of an anhydrous acid. And it is dissociation that determines the physical properties of this substance.
Physical properties of nitric acid
An ionized compound, even if partially ionized, is difficult to convert into a gas. Thus, the boiling point should be quite high, but with such a small molecular weight (and due to high mobility), the melting point should not be high. Consequently, the state of aggregation at 20 °C is liquid.
As for solubility in water, like many other polar liquids, nitric acid is easily mixed with water in any ratio.
Pure nitric acid is colorless and odorless. However, due to decomposition into oxygen and nitrogen oxide (IV), which also dissolves in it, we can say that ordinary concentrated nitric acid has a yellow-brown color and a pungent odor characteristic of NO 2.
Let's see how the structure of the nitric acid molecule affects its chemical properties.
Chemical properties of nitric acid
The main thing we should note is that the presence of a higher oxidation state of the nitrogen atom limits the properties of nitric acid; it does not react with oxidizing agents. But with reducing agents, primarily with metals, it reacts in an unconventional and varied manner.
1. Reactions with metals. Nitric acid reacts with metals as a strong oxidizing agent even in dilute solutions (unlike sulfuric acid, which exhibits its oxidizing properties only in concentrated form). Metal nitrate is usually formed, but instead of hydrogen, gaseous nitrogen compounds are released: NO 2, NO, N 2 O, N 2 or ammonia, which in an acidic environment immediately turns into ammonium ion. In principle, when a metal reacts with nitric acid, this entire “bouquet” of gases is formed, but depending on the metal and the concentration of the acid, certain components will prevail.
Thus, in laboratory conditions, nitrogen(II) oxide is usually obtained by reacting copper shavings with nitric acid with a density of 1.2 g/cm3, i.e., when copper is treated with dilute acid, this oxide clearly prevails in the gaseous reaction products:
But when nitric acid of the same density (and therefore concentration) reacts with iron, the content of nitrogen oxide (II) in the mixture is only 40% - less than half, and the remaining 60% is evenly distributed between ammonium nitrate, nitrogen, nitric oxide (I ) and nitric oxide (IV) (Fig. 8).
An interesting and vitally important fact should be noted that neither iron nor aluminum react with 100% nitric acid (hence, it can be stored and transported in tanks and other containers made of these metals). The fact is that these metals are covered with durable films of oxides that are insoluble in pure acid. For acidic properties to manifest, the acid must be noticeably dissociated, and this in turn requires water.
2. Reactions with non-metals. Nitric acid does not react with oxygen and ozone.
3. There is no reaction with water. Water only promotes the dissociation of the acid.
4. Reactions with acids. Nitric acid does not react with other acids through exchange or compound reactions. However, it is quite capable of reacting as a strong oxidizing agent. In a mixture of concentrated nitric and hydrochloric acids, reversible reactions occur, the essence of which can be summarized by the equation:
The resulting atomic chlorine is very active and easily takes electrons from metal atoms, and the chloride ion available “nearby” forms stable complex ions with the resulting metal ions. All this allows even gold to be transferred into solution. Due to the fact that gold is the “king of metals,” a mixture of concentrated nitric and hydrochloric acids is called aqua regia.
Concentrated sulfuric acid, as a strong water-removing agent, promotes the decomposition of nitric acid into nitric oxide (IV) and oxygen.
5. Reactions with bases and basic oxides. Nitric acid is one of the strong inorganic acids and naturally reacts with alkalis. It also reacts with insoluble hydroxides and basic oxides. These reactions are also facilitated by the fact that all nitric acid salts have good solubility in water, therefore, the reaction products will not interfere with its progress.
Physical properties of NO, NO 2 and HNO 3 compounds in numbers
Nitric oxide(II) NO. Molar mass 30 g/mol. The melting point is –164 °C, the boiling point is –154 °C. The density of gaseous NO under normal conditions (0 °C, 1 atm) is 1.3402 g/l. Solubility at atmospheric pressure and 20 °C is 4.7 ml of NO gas per 100 g of water.
Nitric oxide(IV)) NO 2 . Molar mass 46 g/mol. Melting point –11 °C, boiling point 21 °C. Density of gaseous NO 2 at n. u. 1.491 g/l. Solubility - provided that this oxide first reacts with water in air and then also dissolves in the resulting nitric acid - can be considered unlimited (up to the formation of 60% HNO 3).
Since nitrogen oxide (IV) actively dimerizes (at 140 °C it is entirely in the form of the NO 2 monomer, but at 40 °C about 30% of the monomer remains, and at 20 °C almost all of it turns into the N 2 O 4 dimer) , then the physical properties relate to a dimer rather than a monomer. This is what can explain the rather high boiling point (N 2 O 4 is a fairly heavy molecule). The degree of dimerization can be judged by color: the monomer is intensely colored, and the dimer is colorless.
Nitric acid HNO3. Molar mass 63 g/mol. Melting point –41.6 °C, boiling point 83 °C. The density of liquid 100% acid is 1.513 g/cm3. Solubility is unlimited, in other words, acid and water are mixed in any ratio. It is worth noting that solutions of nitric acid boil at temperatures higher than the boiling points of pure water and acid. At the maximum temperature (122 °C), a 68.4% solution boils, while the percentage composition of the solution and steam is the same.
Mixtures of substances for which the composition of the vapor at boiling corresponds to the composition of the liquid are called azeotropic or non-separately boiling. (The word “azeotrope” comes from the Greek - boil, - change, - negative prefix.) Lower acid concentrations are characterized by an increase in the amount of water in the vapor compared to the solution, which leads to concentration of the solution. At higher concentrations, on the contrary, the vapor composition is enriched with acid.
Chemical properties of nitrogen compounds (supplement)
Like any other substances containing an atom with an intermediate oxidation state, oxides of nitrogen (II) and (IV), unlike nitric acid, can act as both oxidizing agents and reducing agents, depending on the reaction partner. However, many of these reactions are “irrelevant” and, accordingly, poorly studied.
Among the “current” reactions, it is worth mentioning the reaction of nitrogen(IV) oxide with sulfur(IV) oxide in the presence of water:
This reaction is relevant because the addition of oxygen to sulfur(IV) oxide occurs only at high temperatures and in the presence of a catalyst, while the addition of oxygen to nitrogen oxide(II) occurs under normal conditions. Thus, nitrogen(IV) oxide simply helps sulfur oxide to add oxygen. This reaction occurs under normal conditions (additional pressure in the mixture and heating are not required).
Nitric oxide(II) also reacts with sulfur(IV) oxide, but under completely different conditions: either at a pressure of 500 atmospheres (!), then sulfur(VI) oxide and nitrogen are obtained, or in an aqueous solution, then sulfuric acid and nitric oxide(I) are obtained ).
Nitric oxide(I). It has a faint pleasant odor and a sweetish taste. Does not react with oxygen, water, solutions of acids and alkalis. It decomposes into elements at temperatures above 500 °C, in other words, it is quite stable.
The structure of the molecule is interesting: a linear molecule O=N=N, in which the central nitrogen atom is tetravalent. It forms two double bonds: one with oxygen according to the typical scheme of creating a covalent bond (two electrons of nitrogen, two electrons of oxygen), the other with a nitrogen atom (which pairs two of its three unpaired electrons and thereby forms an empty orbital), one of the bonds is covalent, the second is donor-acceptor (Fig. 9).
 |
Rice. 9. |
Nitric oxide (III). Consists of NO and NO 2, which have paired their unpaired electrons. It begins to decompose into the corresponding gases already upon melting (–101 ° C).
Nitric oxide(V). Consists of two NO 2 groups connected through oxygen. A somewhat more stable compound than nitrogen(III) oxide, it begins to decompose at room temperature. Some of the bonds in it are, naturally, donor-acceptor. And no “pentavalent nitrogens”.
It should be added to the chemical properties of nitric acid that it reacts well with non-metals that it can oxidize. Thus, concentrated nitric acid reacts with sulfur, phosphorus, and coal, forming sulfuric, phosphoric and carbonic acids, respectively.
The reactions of nitric acid with organic substances are interesting and important. For example, when three hydrogen atoms in toluene are replaced by NO 2 groups, trinitrotoluene (or simply tol) is formed, an explosive substance.
Environmental properties of nitrogen oxides and nitric acid
Nitric oxide(I) relatively inert and therefore “ecologically neutral”. However, it has a narcotic effect on humans, ranging from simple fun (for which it was nicknamed “laughing gas”) and ending with deep sleep, which has found its application in medicine. Interestingly, it is harmless, and for medical anesthesia a mixture of nitric oxide (I) with oxygen is used in the same ratio as the ratio of nitrogen and oxygen in the air. The narcotic effect is removed immediately after you stop inhaling this gas.
The other two stable nitrogen oxides easily transform into one another, then into acids, and then into anions and. Thus, these substances are natural mineral fertilizers, however, if they are found in natural quantities. In “unnatural” quantities, these gases rarely enter the atmosphere alone. As a rule, a whole “bouquet” of toxic compounds is formed that act in a complex manner.
For example, just one nitrogen fertilizer plant emits into the air, in addition to nitrogen oxides, nitric acid, ammonia and fertilizer dust, sulfur oxides, fluorine compounds, and some organic compounds. Scientists are determining the resistance of various grasses, bushes and trees to such “bouquets”. It is already known that, unfortunately, spruce and pine are unstable and die quickly, but white acacia, Canadian poplar, willows and some other plants can exist in such conditions, moreover, they help remove these substances from the air.
Severe poisoning with nitrogen oxides can occur mainly during accidents at relevant industries. The body's response will be different due to the differences in the properties of these gases. “Caustic” NO 2 primarily affects the mucous membranes of the nasopharynx and eyes, causing pulmonary edema; NO, as a poorly soluble and non-caustic substance in water, passes through the lungs and enters the blood, causing disturbances in the central and peripheral nervous systems. Both oxides react with hemoglobin in the blood, although in different ways, but with the same result - hemoglobin stops carrying oxygen.
The environmental properties of nitric acid consist of two “halves”. As a strong acid, it has a destructive effect not only on living tissue (human skin, plant leaves), but also on the soil, which is quite important - acidic (due to the presence of nitrogen and sulfur oxides) rains, alas, are not uncommon. When acid gets on the skin, a chemical burn occurs, which is more painful and takes much longer to heal than a thermal burn. These were the main environmental properties hydrogen cation.
Let's move on to studying anion. When a strong acid acts, it is the acidic properties that come to the fore, so it is better to consider the properties of the anion using the example of salts.
Interaction of nitrate ion with fauna and flora. The fact is that the nitrate ion is an integral part of the nitrogen cycle in nature; it is always present in it. Under normal conditions and in dilute solutions, it is stable, exhibits weak oxidizing properties, and does not precipitate metal cations, thereby facilitating the transport of these ions with the solution in soil, plants, etc.
Nitrate ion becomes toxic only in large quantities, disturbing the balance of other substances. For example, with an excess of nitrates in plants, the amount of ascorbic acid decreases. (It is worth recalling that a living organism is so finely organized that any substance in large quantities upsets the balance and, therefore, becomes poisonous.)
Plants and bacteria use nitrates to build proteins and other essential organic compounds. To do this, you need to convert the nitrate ion into an ammonium ion. This reaction is catalyzed by enzymes containing metal ions (copper, iron, manganese, etc.). Due to the much greater toxicity of ammonia and ammonium ion in plants, the reverse reaction of converting ammonium ion into nitrate is well developed.
Animals do not know how to build all the organic compounds they need from inorganic ones - they lack the appropriate enzymes. However, microorganisms living in the stomach and intestines possess these enzymes and can convert nitrate ion into nitrite ion. It is the nitrite ion that acts as a poisoner, converting the iron in hemoglobin from Fe 2+ to Fe 3+.
The compound containing Fe 3+ and called methemoglobin binds air oxygen too tightly, therefore, cannot release it to the tissues. As a result, the body suffers from a lack of oxygen, and disturbances in the functioning of the brain, heart and other organs occur.
Typically, the nitrite ion is formed not in the stomach, but in the intestines and does not have time to pass into the blood and cause all this destruction. Therefore, nitrate poisoning is quite rare. There is, however, another danger: in our body there are many substances in which the hydrogen atoms of ammonia are replaced by organic radicals. Such compounds are called amines. When amines react with nitrite ions, nitrosamines are formed - carcinogenic substances:
They act on the liver and promote the formation of tumors in the lungs and kidneys. Interestingly, the active inhibitor of the reaction of nitrosamine formation is ascorbic acid, which has long been familiar to us.
O.R.VALEDINSKAYA
(MSU, Moscow)
