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As a result of studying this topic, you will learn:
As a result of studying this topic, you will learn:
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5.1. covalent bond
A chemical bond is formed when two or more more atoms, if as a result of their interaction there is a decrease in the total energy of the system. The most stable electronic configurations of the outer electron shells of atoms are configurations of noble gas atoms, consisting of two or eight electrons. The outer electron shells of atoms of other elements contain from one to seven electrons, i.e. are incomplete. When a molecule is formed, atoms tend to acquire a stable two-electron or eight-electron shell. The valence electrons of atoms take part in the formation of a chemical bond.
A covalent bond is a chemical bond between two atoms, which is formed by electron pairs belonging simultaneously to these two atoms.
There are two mechanisms for the formation of a covalent bond: exchange and donor-acceptor.
5.1.1. Exchange mechanism for the formation of a covalent bond
exchange mechanism The formation of a covalent bond is realized due to the overlapping of electron clouds of electrons belonging to different atoms. For example, when two hydrogen atoms approach each other, the 1s electron orbitals overlap. As a result, a common pair of electrons appears, simultaneously belonging to both atoms. In this case, the chemical bond is formed by electrons having antiparallel spins, Fig. 5.1.
Rice. 5.1. Formation of a hydrogen molecule from two H atoms
5.1.2. Donor-acceptor mechanism of covalent bond formation
With the donor-acceptor mechanism for the formation of a covalent bond, the bond is also formed with the help of electron pairs. However, in this case, one atom (donor) provides its electron pair, and the other atom (acceptor) participates in the formation of the bond with its free orbital. An example of the implementation of a donor-acceptor bond is the formation of an ammonium ion NH 4 + during the interaction of ammonia NH 3 with a hydrogen cation H + .
In the NH 3 molecule, three electron pairs form three N - H bonds, the fourth electron pair belonging to the nitrogen atom is unshared. This electron pair can give a bond to the hydrogen ion, which has a free orbital. The result is an ammonium ion NH 4 + , fig. 5.2.
Rice. 5.2. Occurrence of a donor-acceptor bond during the formation of an ammonium ion
It should be noted that the four N – H covalent bonds existing in the NH 4 + ion are equivalent. In the ammonium ion, it is impossible to isolate the bond formed by the donor-acceptor mechanism.
5.1.3. Polar and non-polar covalent bond
If a covalent bond is formed by identical atoms, then the electron pair is located at the same distance between the nuclei of these atoms. Such a covalent bond is called non-polar. An example of molecules with a non-polar covalent bond are H 2, Cl 2, O 2, N 2, etc.
In the case of a polar covalent bond, the shared electron pair is shifted towards the atom with the higher electronegativity. This type of bond is realized in molecules formed by different atoms. The covalent polar bond takes place in the molecules of HCl, HBr, CO, NO, etc. For example, the formation of a polar covalent bond in the HCl molecule can be represented by the scheme, fig. 5.3:
Rice. 5.3. Formation of a covalent polar bond in the HC1 molecule
In the molecule under consideration, the electron pair is shifted to the chlorine atom, since its electronegativity (2.83) is greater than the electronegativity of the hydrogen atom (2.1).
5.1.4. Dipole moment and structure of molecules
The measure of bond polarity is its dipole moment μ:
μ = e l,
Where e is the charge of an electron, l is the distance between the centers of positive and negative charges.
The dipole moment is vector quantity. The concepts of "bond dipole moment" and "dipole moment of a molecule" coincide only for diatomic molecules. The dipole moment of a molecule is equal to the vector sum of the dipole moments of all bonds. Thus, the dipole moment of a polyatomic molecule depends on its structure.
In a linear CO 2 molecule, for example, each of the C–O bonds is polar. However, the CO 2 molecule is generally non-polar, since the dipole moments of the bonds compensate each other (Fig. 5.4). The dipole moment of a carbon dioxide molecule is m = 0.
In the corner H 2 O molecule, the polar H–O bonds are located at an angle of 104.5 o. The vector sum of the dipole moments of two H–O bonds is expressed by the diagonal of the parallelogram (Fig. 5.4). As a result, the dipole moment of the water molecule m is not equal to zero.
Rice. 5.4. Dipole moments of CO 2 and H 2 O molecules
5.1.5. Valency of elements in compounds with a covalent bond
The valency of atoms is determined by the number of unpaired electrons participating in the formation of common electron pairs with electrons of other atoms. Having one unpaired electron on the outer electron layer, the halogen atoms in the F 2, HCl, PBr 3 and CCl 4 molecules are monovalent. Elements of the oxygen subgroup contain two unpaired electrons on the outer layer, so in compounds such as O 2, H 2 O, H 2 S and SCl 2 they are divalent.
Since, in addition to the usual covalent bonds, a bond can be formed in molecules by a donor-acceptor mechanism, the valence of atoms also depends on the presence of lone electron pairs and free electron orbitals in them. A quantitative measure of valence is the number of chemical bonds by which a given atom is connected to other atoms.
The maximum valence of elements, as a rule, cannot exceed the number of the group in which they are located. The exception is the elements of the side subgroup of the first group Cu, Ag, Au, whose valency in compounds is greater than one. The electrons of the outer layers primarily belong to the valence ones, however, for the elements of the secondary subgroups, the electrons of the penultimate (anterior) layers also take part in the formation of a chemical bond.
5.1.6. Valency of elements in normal and excited states
Majority Valence chemical elements depends on whether these elements are in the normal or excited state. Electronic configuration of the Li atom: 1s 2 2s 1. The lithium atom at the outer level has one unpaired electron, i.e. lithium is monovalent. A very large expenditure of energy is required, associated with the transition of a 1s electron to a 2p orbital, in order to obtain trivalent lithium. This energy expenditure is so great that it is not compensated by the energy released during the formation of chemical bonds. In this regard, there are no compounds of trivalent lithium.
The configuration of the outer electron layer of the elements of the subgroup of beryllium ns 2 . This means that on the outer electron layer of these elements, there are two electrons with opposite spins in the ns cell orbital. The elements of the beryllium subgroup do not contain unpaired electrons, so their valency in the normal state is zero. In the excited state, the electronic configuration of the elements of the beryllium subgroup is ns 1 nр 1, i.e. elements form compounds in which they are divalent.
Valence possibilities of the boron atom
Consider the electronic configuration of the boron atom in the ground state: 1s 2 2s 2 2р 1 . The boron atom in the ground state contains one unpaired electron (Fig. 5.5), i.e. he is univalent. However, boron is not characterized by the formation of compounds in which it is monovalent. When a boron atom is excited, a transition of one 2s-electron to a 2p-orbital occurs (Fig. 5.5). The boron atom in an excited state has 3 unpaired electrons and can form compounds in which its valency is three.
Rice. 5.5. Valence states of the boron atom in the normal and excited states
The energy spent on the transition of an atom to an excited state within one energy level, as a rule, is compensated in excess by the energy released during the formation of additional bonds.
Due to the presence of one free 2p orbital in the boron atom, boron in compounds can form a fourth covalent bond, acting as an electron pair acceptor. Figure 5.6 shows how the BF molecule interacts with the F ion - , as a result of which an ion - is formed, in which boron forms four covalent bonds.
Rice. 5.6. Donor-acceptor mechanism for the formation of the fourth covalent bond at the boron atom
Valence possibilities of the nitrogen atom
Consider the electronic structure of the nitrogen atom (Fig. 5.7).
Rice. 5.7. The distribution of electrons in the orbitals of the nitrogen atom
From the presented diagram it can be seen that nitrogen has three unpaired electrons, it can form three chemical bonds and its valency is three. The transition of the nitrogen atom to an excited state is impossible, since the second energy level does not contain d-orbitals. At the same time, the nitrogen atom can provide an unshared electron pair of outer electrons 2s 2 to an atom that has a free orbital (acceptor). As a result, a fourth chemical bond of the nitrogen atom arises, as is the case, for example, in the ammonium ion (Fig. 5.2). Thus, the maximum covalence (the number of formed covalent bonds) of the nitrogen atom is four. In its compounds, nitrogen, unlike other elements of the fifth group, cannot be pentavalent.
Valence possibilities of phosphorus, sulfur and halogen atoms
Unlike nitrogen, oxygen, and fluorine atoms, phosphorus, sulfur, and chlorine atoms in the third period have free 3d cells, to which electrons can transfer. When a phosphorus atom is excited (Fig. 5.8), it has 5 unpaired electrons on its outer electron layer. As a result, in compounds, the phosphorus atom can be not only tri-, but also pentavalent.
Rice. 5.8. Distribution of valence electrons in orbits for a phosphorus atom in an excited state
In an excited state, sulfur, in addition to a valency of two, also exhibits a valence of four and six. In this case, the depairing of 3p and 3s electrons occurs sequentially (Fig. 5.9).
Rice. 5.9. Valence possibilities of the sulfur atom in an excited state
In the excited state, for all elements of the main subgroup of group V, except for fluorine, sequential depairing of first p- and then s-electron pairs is possible. As a result, these elements become tri-, penta-, and heptavalent (Fig. 5.10).
Rice. 5.10. Valence possibilities of chlorine, bromine and iodine atoms in an excited state
5.1.7. Length, energy and direction of a covalent bond
A covalent bond, as a rule, is formed between the atoms of non-metals. The main characteristics of a covalent bond are length, energy, and directionality.
Covalent bond length
The bond length is the distance between the nuclei of the atoms that form this bond. It is determined by experimental physical methods. The bond length can be estimated using the additivity rule, according to which the bond length in the AB molecule is approximately equal to half the sum of the bond lengths in the A 2 and B 2 molecules:
.
From top to bottom in the subgroups of the periodic system of elements, the length of the chemical bond increases, since the radii of atoms increase in this direction (Table 5.1). As the bond multiplicity increases, its length decreases.
Table 5.1.
The length of some chemical bonds
chemical bond |
Communication length, pm |
chemical bond |
Communication length, pm |
C - C |
|||
The measure of bond strength is the bond energy. Bond energy is determined by the energy required to break the bond and remove the atoms that form this bond to an infinite distance from each other. The covalent bond is very strong. Its energy ranges from several tens to several hundreds of kJ/mol. For an IСl 3 molecule, for example, Ebonds ≈40, and for N 2 and CO molecules, Ebonds ≈1000 kJ/mol.
From top to bottom in the subgroups of the periodic system of elements, the energy of a chemical bond decreases, since the bond length increases in this direction (Table 5.1). With an increase in the multiplicity of the connection, its energy increases (Table 5.2).
Table 5.2.
The energies of some chemical bonds
chemical bond |
bond energy, |
chemical bond |
bond energy, |
C - C |
|||
Saturation and directionality of a covalent bond
The most important properties of a covalent bond are its saturation and directionality. Saturation can be defined as the ability of atoms to form a limited number of covalent bonds. So a carbon atom can form only four covalent bonds, and an oxygen atom can form two. The maximum number of ordinary covalent bonds that an atom can form (excluding bonds formed by the donor-acceptor mechanism) is equal to the number of unpaired electrons.
Covalent bonds have a spatial orientation, since the overlap of orbitals during the formation of a single bond occurs along the line connecting the nuclei of atoms. The spatial arrangement of the electron orbitals of a molecule determines its geometry. The angles between chemical bonds are called bond angles.
The saturation and directionality of a covalent bond distinguishes this bond from an ionic bond, which, unlike a covalent bond, is unsaturated and non-directional.
Spatial structure of H 2 O and NH 3 molecules
Let us consider the orientation of a covalent bond using the example of H 2 O and NH 3 molecules.
The H 2 O molecule is formed from an oxygen atom and two hydrogen atoms. The oxygen atom has two unpaired p-electrons that occupy two orbitals located at right angles to each other. Hydrogen atoms have unpaired 1s electrons. The angle between the bonds formed by the p-electrons should be close to the angle between the orbitals of the p-electrons. Experimentally, however, it was found that the angle between the O–H bonds in a water molecule is 104.50. The increase in the angle compared to the angle of 90 o can be explained by the repulsive forces that act between the hydrogen atoms, fig. 5.11. Thus, the H 2 O molecule has an angular shape.
Three unpaired p-electrons of the nitrogen atom participate in the formation of the NH 3 molecule, the orbitals of which are located in three mutually perpendicular directions. Therefore, the three N–H bonds must be at angles to each other close to 90° (Fig. 5.11). The experimental value of the angle between bonds in the NH 3 molecule is 107.3°. The difference in the values of the angles between the bonds from the theoretical values is due, as in the case of the water molecule, to the mutual repulsion of the hydrogen atoms. In addition, the presented schemes do not take into account the possibility of participation of two electrons in 2s orbitals in the formation of chemical bonds.
Rice. 5.11. Overlapping of electronic orbitals during the formation of chemical bonds in H 2 O (a) and NH 3 (b) molecules
Consider the formation of the BeCl 2 molecule. An atom of beryllium in an excited state has two unpaired electrons: 2s and 2p. It can be assumed that the beryllium atom should form two bonds: one bond formed by the s-electron and one bond formed by the p-electron. These bonds must have different energies and different lengths. The BeCl 2 molecule in this case should not be linear, but angular. Experience, however, shows that the BeCl 2 molecule has a linear structure and both chemical bonds in it are equivalent. A similar situation is observed when considering the structure of BCl 3 and CCl 4 molecules – all bonds in these molecules are equivalent. The BC1 3 molecule has a planar structure, CC1 4 is tetrahedral.
To explain the structure of molecules such as BeCl 2, BCl 3 and CCl 4, Pauling and Slater(USA) introduced the concept of hybridization of atomic orbitals. They proposed to replace several atomic orbitals, not very different in their energy, with the same number of equivalent orbitals, called hybrid ones. These hybrid orbitals are made up of atomic orbitals as a result of their linear combination.
According to L. Pauling, when chemical bonds are formed by an atom that has electrons of various types in one layer and, therefore, not very different in energy (for example, s and p), it is possible to change the configuration of orbitals of various types, in which they are aligned in shape and energy . As a result, hybrid orbitals are formed, which have an asymmetric shape and are strongly elongated on one side of the nucleus. It is important to emphasize that the hybridization model is used in the case when electrons of different types participate in the formation of bonds, for example, s and p.
5.1.8.2. Various types of hybridization of atomic orbitals
sp hybridization
Hybridization of one s- and one R- orbitals ( sp- hybridization) realized, for example, in the formation of beryllium chloride. As shown above, in the excited state, the Be atom has two unpaired electrons, one of which occupies the 2s orbital and the other, the 2p orbital. When a chemical bond is formed, these two different orbitals are transformed into two identical hybrid orbitals directed at an angle of 180 ° to each other (Fig. 5.12). The linear arrangement of two hybrid orbitals corresponds to their minimum repulsion from each other. As a result, the BeCl 2 molecule has a linear structure - all three atoms are located on the same line.
Rice. 5.12. Scheme of overlapping electron orbitals during the formation of the BeCl 2 molecule
The structure of the acetylene molecule; sigma and pi bonds
Consider the scheme of overlapping electron orbitals in the formation of an acetylene molecule. In the acetylene molecule, each carbon atom is in the sp hybrid state. Two sp-hybrid orbitals are located at an angle of 1800 to each other; they form one σ-bond between carbon atoms and two σ-bonds with hydrogen atoms (Fig. 5.13).
Rice. 5.13. Scheme of the formation of s-bonds in the acetylene molecule
A σ-bond is a bond formed as a result of the overlap of electron orbitals along the line connecting the nuclei of atoms.
Each carbon atom in the acetylene molecule contains two more p-electrons, which do not take part in the formation of σ-bonds. The electron clouds of these electrons are located in mutually perpendicular planes and, overlapping with each other, form two more π bonds between carbon atoms due to the lateral overlap of non-hybrid R-clouds (Fig. 5.14).
A π bond is a covalent chemical bond formed as a result of an increase in electron density on either side of a line connecting the nuclei of atoms.
Rice. 5.14. Scheme of the formation of σ - and π -bonds in the acetylene molecule.
Thus, in an acetylene molecule, a triple bond is formed between carbon atoms, which consists of one σ bond and two π bonds; σ -bonds are stronger than π-bonds.
sp2 hybridization
The structure of the BCl 3 molecule can be explained in terms of sp 2- hybridization. The boron atom in the excited state contains one s-electron and two p-electrons on the outer electron layer, i.e. three unpaired electrons. These three electron clouds can be converted into three equivalent hybrid orbitals. The minimum repulsion of three hybrid orbitals from each other corresponds to their location in the same plane at an angle of 120 o to each other (Fig. 5.15). Thus, the BCl 3 molecule has a planar shape.
Rice. 5.15. The planar structure of the BCl 3 molecule
sp 3 - hybridization
The valence orbitals of the carbon atom (s, p x, p y, p z) can be converted into four equivalent hybrid orbitals, which are located in space at an angle of 109.5 o to each other and are directed to the vertices of the tetrahedron, in the center of which is the nucleus of the carbon atom (Fig. 5.16).
Rice. 5.16. Tetrahedral structure of the methane molecule
5.1.8.3. Hybridization involving lone electron pairs
The hybridization model can be used to explain the structure of molecules in which, in addition to binding, there are also unshared electron pairs. In water and ammonia molecules total number electron pairs of the central atom (O and N) is four. In this case, the water molecule has two, and the ammonia molecule has one unshared electron pair. The formation of chemical bonds in these molecules can be explained by assuming that lone electron pairs can also fill hybrid orbitals. Unshared electron pairs occupy much more space in space than bonding pairs. As a result of the repulsion that occurs between lone and bonding electron pairs, the bond angles in water and ammonia molecules decrease, which turn out to be less than 109.5 o.
Rice. 5.17. sp 3 - hybridization involving lone electron pairs in H 2 O (A) and NH 3 (B) molecules
5.1.8.4. Establishment of the type of hybridization and determination of the structure of molecules
To establish the type of hybridization, and, consequently, the structure of molecules, the following rules must be used.
1. The type of hybridization of the central atom, which does not contain unshared electron pairs, is determined by the number of sigma bonds. If there are two such bonds, sp-hybridization takes place, three - sp 2 -hybridization, four - sp 3 -hybridization. Unshared electron pairs (in the absence of bonds formed by the donor-acceptor mechanism) are absent in molecules formed by atoms of beryllium, boron, carbon, silicon, i.e. the elements of the main subgroups II - IV groups.
2. If the central atom contains unshared electron pairs, then the number of hybrid orbitals and the type of hybridization are determined by the sum of the number of sigma bonds and the number of unshared electron pairs. Hybridization involving unshared electron pairs takes place in molecules formed by nitrogen, phosphorus, oxygen, and sulfur atoms, i.e. elements of the main subgroups of groups V and VI.
3. The geometric shape of the molecules is determined by the type of hybridization of the central atom (Table 5.3).
Table 5.3.
Valence angles geometric shape molecules depending on the number of hybrid orbitals and the type of hybridization of the central atom
5.2. Ionic bond
Ionic bonding is carried out by electrostatic attraction between oppositely charged ions. These ions are formed as a result of the transfer of electrons from one atom to another. An ionic bond is formed between atoms that have large differences in electronegativity (usually greater than 1.7 on the Pauling scale), for example, between atoms alkali metals and halogens.
Let us consider the appearance of an ionic bond using the example of the formation of NaCl. From the electronic formulas of the atoms Na 1s 2 2s 2 2p 6 3s 1 and Cl 1s 2 2s 2 2p 6 3s 2 3p 5, it is easier to give one electron to the sodium atom than to attach seven, and it is easier for the chlorine atom to attach one, than give seven. In chemical reactions, the sodium atom donates one electron, and the chlorine atom accepts it. As a result, the electron shells of sodium and chlorine atoms turn into stable electron shells of noble gases (the electronic configuration of the sodium cation is Na + 1s 2 2s 2 2p 6, and the electronic configuration of the chlorine anion Cl is 1s 2 2s 2 2p 6 3s 2 3p 6). The electrostatic interaction of ions leads to the formation of the NaCl molecule.
Main characteristics of ionic bond and properties of ionic compounds
1. An ionic bond is a strong chemical bond. The energy of this bond is about 300 – 700 kJ/mol.
2. Unlike a covalent bond, an ionic bond is non-directional, since an ion can attract ions of the opposite sign to itself in any direction.
3. Unlike a covalent bond, an ionic bond is unsaturated, since the interaction of ions of the opposite sign does not lead to complete mutual compensation of their force fields.
4. In the process of formation of molecules with an ionic bond, there is no complete transfer of electrons, therefore, a 100% ionic bond does not exist in nature. In the NaCl molecule, the chemical bond is only 80% ionic.
5. Ionic compounds are crystalline solids with high melting and boiling points.
6. Most ionic compounds dissolve in water. Solutions and melts of ionic compounds conduct electricity.
5.3. metal connection
Metal atoms at the outer energy level contain a small number of valence electrons. Since the ionization energy of metal atoms is low, valence electrons are weakly retained in these atoms. As a result, positively charged ions and free electrons appear in the crystal lattice of metals. In this case, the metal cations are located at the nodes of their crystal lattice, and the electrons move freely in the field of positive centers, forming the so-called "electron gas". The presence of a negatively charged electron between two cations leads to the fact that each cation interacts with this electron. Thus, a metallic bond is a bond between positive ions in metal crystals, which is carried out by the attraction of electrons moving freely throughout the crystal.
Since the valence electrons in the metal are evenly distributed throughout the crystal, the metallic bond, like the ionic one, is an undirected bond. Unlike a covalent bond, a metallic bond is an unsaturated bond. From a covalent bond metallic bond differs also in durability. The energy of a metallic bond is about three to four times less than the energy of a covalent bond.
Due to the high mobility of the electron gas, metals are characterized by high electrical and thermal conductivity.
5.4. hydrogen bond
In the molecules of compounds HF, H 2 O, NH 3, there are hydrogen bonds with a strongly electronegative element (H–F, H–O, H–N). Between the molecules of such compounds can be formed intermolecular hydrogen bonds. In some organic molecules containing H–O, H–N bonds, intramolecular hydrogen bonds.
The mechanism of hydrogen bond formation is partly electrostatic, partly donor-acceptor. In this case, the atom of a strongly electronegative element (F, O, N) acts as an electron pair donor, and the hydrogen atoms connected to these atoms act as an acceptor. As with covalent bonds, hydrogen bonds are characterized by orientation in space and saturability.
The hydrogen bond is usually denoted by dots: H ··· F. The hydrogen bond is more pronounced, the greater the electronegativity of the partner atom and the smaller its size. It is characteristic primarily for fluorine compounds, as well as oxygen, to a lesser extent nitrogen, to an even lesser extent for chlorine and sulfur. Accordingly, the energy of the hydrogen bond also changes (Table 5.4).
Table 5.4.
Average values of hydrogen bond energies
Intermolecular and intramolecular hydrogen bonding
Thanks to hydrogen bonds, molecules are combined into dimers and more complex associates. For example, the formation of a formic acid dimer can be represented by the following scheme (Fig. 5.18).
Rice. 5.18. Formation of intermolecular hydrogen bonds in formic acid
Long chains of associates (H 2 O) n can appear in water (Fig. 5.19).
Rice. 5.19. Formation of a chain of associates in liquid water due to intermolecular hydrogen bonds
Each H 2 O molecule can form four hydrogen bonds, while an HF molecule can form only two.
Hydrogen bonds can occur both between different molecules (intermolecular hydrogen bond) and within a molecule (intramolecular hydrogen bond). Examples of intramolecular bond formation for some organic matter are presented in fig. 5.20.
Rice. 5.20. Formation of an intramolecular hydrogen bond in the molecules of various organic compounds
The effect of hydrogen bonding on the properties of substances
The most convenient indicator of the existence of an intermolecular hydrogen bond is the boiling point of a substance. The higher boiling point of water (100 o C compared to hydrogen compounds elements of the oxygen subgroup (H 2 S, H 2 Se, H 2 Te) is explained by the presence of hydrogen bonds: it is necessary to spend additional energy on the destruction of intermolecular hydrogen bonds in water.
The hydrogen bond can significantly affect the structure and properties of substances. The existence of intermolecular hydrogen bonds increases the melting and boiling points of substances. The presence of an intramolecular hydrogen bond leads to the fact that the molecule of deoxyribonucleic acid (DNA) is folded into a double helix in water.
The hydrogen bond also plays important role in dissolution processes, since solubility also depends on the ability of the compound to form hydrogen bonds with the solvent. As a result, substances containing OH groups such as sugar, glucose, alcohols, carboxylic acids, as a rule, are highly soluble in water.
5.5. Types of crystal lattices
Solids, as a rule, have a crystalline structure. The particles that make up crystals (atoms, ions or molecules) are located at strictly defined points in space, forming a crystal lattice. The crystal lattice consists of elementary cells that retain the structural features characteristic of this lattice. The points where the particles are located are called lattice nodes. Depending on the type of particles located at the lattice sites and on the nature of the connection between them, 4 types of crystal lattices are distinguished.
5.5.1. Atomic crystal lattice
At the nodes of atomic crystal lattices there are atoms interconnected by covalent bonds. Substances having an atomic lattice include diamond, silicon, carbides, silicides, etc. In the structure of an atomic crystal, it is impossible to single out individual molecules; the entire crystal is considered as one giant molecule. The structure of diamond is shown in fig. 5.21. A diamond is made up of carbon atoms, each bonded to four neighboring atoms. Due to the fact that covalent bonds are strong, all substances having atomic lattices are refractory, solid and low volatile. They are slightly soluble in water.
Rice. 5.21. Diamond crystal lattice
5.5.2. Molecular crystal lattice
Molecules are located at the nodes of molecular crystal lattices, interconnected by weak intermolecular forces. Therefore, substances with a molecular lattice have low hardness, they are fusible, are characterized by significant volatility, are slightly soluble in water, and their solutions, as a rule, do not conduct electric current. A lot of substances with a molecular crystal lattice are known. These are solid hydrogen, chlorine, carbon monoxide (IV) and other substances that are in a gaseous state at ordinary temperatures. Most crystalline organic compounds have a molecular lattice.
5.5.3. Ionic crystal lattice
Crystal lattices, at the nodes of which ions are located, are called ionic. They are formed by substances with an ionic bond, for example, alkali metal halides. In ionic crystals, individual molecules cannot be distinguished; the entire crystal can be considered as one macromolecule. The bonds between ions are strong, so substances with an ionic lattice have low volatility, high melting and boiling points. The crystal lattice of sodium chloride is shown in fig. 5.22.
Rice. 5.22. Crystal lattice of sodium chloride
In this figure, light balls are Na + ions, dark balls are Cl - ions. On the left in fig. 5.22 shows the unit cell of NaCI.
5.5.4. metal crystal lattice
Metals in the solid state form metallic crystal lattices. At the nodes of such lattices there are positive metal ions, and valence electrons move freely between them. The electrons electrostatically attract the cations, thereby giving stability to the metal lattice. Such a structure of the lattice determines the high thermal conductivity, electrical conductivity and plasticity of metals - mechanical deformation does not break the bonds and destroy the crystal, since the ions that make it up seem to float in a cloud of electron gas. On fig. 5.23 shows the crystal lattice of sodium.
Rice. 5.23. The crystal lattice of sodium
.
You know that atoms can combine with each other to form both simple and complex substances. In this case, various types of chemical bonds are formed: ionic, covalent (non-polar and polar), metallic and hydrogen. One of the most essential properties of the atoms of elements, which determine what kind of bond is formed between them - ionic or covalent, - is the electronegativity, i.e. the ability of atoms in a compound to attract electrons to itself.
conditional quantification electronegativity scale gives a scale of relative electronegativity.
In periods, there is a general tendency for the growth of the electronegativity of the elements, and in groups - their decline. Electronegativity elements are arranged in a row, on the basis of which it is possible to compare the electronegativity of elements in different periods.
The type of chemical bond depends on how large the difference in the electronegativity values of the connecting atoms of the elements is. The more the atoms of the elements forming the bond differ in electronegativity, the more polar the chemical bond is. Conduct sharp border between types of chemical bonds is impossible. In most compounds, the type of chemical bond is intermediate; for example, a highly polar covalent chemical bond is close to an ionic bond. Depending on which of the limiting cases is closer in nature to the chemical bond, it is referred to as either an ionic or a covalent polar bond.
Ionic bond.An ionic bond is formed by the interaction of atoms that differ sharply from each other in electronegativity. For example, typical metals lithium (Li), sodium (Na), potassium (K), calcium (Ca), strontium (Sr), barium (Ba) form an ionic bond with typical non-metals, mainly halogens.
In addition to alkali metal halides, ionic bonds are also formed in compounds such as alkalis and salts. For example, in sodium hydroxide (NaOH) and sodium sulfate (Na 2 SO 4) ionic bonds exist only between sodium and oxygen atoms (the rest of the bonds are covalent polar).
Covalent non-polar bond.When atoms interact with the same electronegativity, molecules are formed with a covalent non-polar bond. Such a bond exists in the molecules of the following simple substances: H 2 , F 2 , Cl 2 , O 2 , N 2 . Chemical bonds in these gases are formed through common electron pairs, i.e. when the corresponding electron clouds overlap, due to the electron-nuclear interaction, which occurs when the atoms approach each other.
Composing electronic formulas substances, it should be remembered that each common electron pair is a conditional image of an increased electron density resulting from the overlap of the corresponding electron clouds.
covalent polar bond.During the interaction of atoms, the values of the electronegativity of which differ, but not sharply, there is a shift of the common electron pair to a more electronegative atom. This is the most common type of chemical bond found in both inorganic and organic compounds.
Covalent bonds fully include those bonds that are formed by the donor-acceptor mechanism, for example, in hydronium and ammonium ions.
Metal connection.
The bond that is formed as a result of the interaction of relatively free electrons with metal ions is called a metallic bond. This type of bond is typical for simple substances - metals.
The essence of the process of formation of a metallic bond is as follows: metal atoms easily give up valence electrons and turn into positively charged ions. Relatively free electrons, detached from the atom, move between positive metal ions. A metallic bond arises between them, i.e., the electrons, as it were, cement the positive ions of the crystal lattice of metals.
Hydrogen bond.
A bond that forms between the hydrogen atoms of one molecule and an atom of a strongly electronegative element(O, N, F) another molecule is called a hydrogen bond.
The question may arise: why exactly does hydrogen form such a specific chemical bond?
This is because the atomic radius of hydrogen is very small. In addition, when a single electron is displaced or completely donated, hydrogen acquires a relatively high positive charge, due to which the hydrogen of one molecule interacts with atoms of electronegative elements that have a partial negative charge that is part of other molecules (HF, H 2 O, NH 3) .
Let's look at some examples. We usually depict the composition of water chemical formula H 2 O. However, this is not entirely accurate. It would be more correct to denote the composition of water by the formula (H 2 O) n, where n \u003d 2.3.4, etc. This is due to the fact that individual water molecules are interconnected through hydrogen bonds.
Hydrogen bonds are usually denoted by dots. It is much weaker than an ionic or covalent bond, but stronger than the usual intermolecular interaction.
The presence of hydrogen bonds explains the increase in the volume of water with decreasing temperature. This is due to the fact that as the temperature decreases, the molecules become stronger and therefore the density of their “packing” decreases.
When studying organic chemistry The following question also arose: why are the boiling points of alcohols much higher than those of the corresponding hydrocarbons? This is explained by the fact that hydrogen bonds are also formed between alcohol molecules.
An increase in the boiling point of alcohols also occurs due to the enlargement of their molecules.
The hydrogen bond is also characteristic of many other organic compounds (phenols, carboxylic acids, etc.). From organic chemistry courses and general biology you know that the presence of a hydrogen bond explains the secondary structure of proteins, the structure of the double helix of DNA, i.e., the phenomenon of complementarity.
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Each atom has a certain number of electrons.
Entering into chemical reactions, atoms donate, acquire, or socialize electrons, reaching the most stable electronic configuration. The configuration with the lowest energy is the most stable (as in noble gas atoms). This pattern is called the "octet rule" (Fig. 1).
Rice. 1.
This rule applies to all connection types. Electronic bonds between atoms allow them to form stable structures, from the simplest crystals to complex biomolecules that eventually form living systems. They differ from crystals in their continuous metabolism. However, many chemical reactions proceed according to the mechanisms electronic transfer, which play an important role in the energy processes in the body.
A chemical bond is a force that holds together two or more atoms, ions, molecules, or any combination of them..
The nature of the chemical bond is universal: it is an electrostatic force of attraction between negatively charged electrons and positively charged nuclei, determined by the configuration of the electrons in the outer shell of atoms. The ability of an atom to form chemical bonds is called valency, or oxidation state. Valence is related to the concept of valence electrons- electrons that form chemical bonds, that is, those located in the most high-energy orbitals. Accordingly, the outer shell of an atom containing these orbitals is called valence shell. At present, it is not enough to indicate the presence of a chemical bond, but it is necessary to clarify its type: ionic, covalent, dipole-dipole, metallic.
The first type of connection isionic connection
In accordance with electronic theory Lewis and Kossel valencies, atoms can achieve a stable electron configuration in two ways: first, by losing electrons, becoming cations, secondly, acquiring them, turning into anions. As a result of electron transfer, due to the electrostatic force of attraction between ions with charges of the opposite sign, a chemical bond is formed, called Kossel " electrovalent(now called ionic).
In this case, anions and cations form a stable electronic configuration with a filled outer electron shell. Typical ionic bonds are formed from cations of T and II groups of the periodic system and anions of non-metallic elements of groups VI and VII (16 and 17 subgroups - respectively, chalcogens And halogens). The bonds in ionic compounds are unsaturated and non-directional, so they retain the possibility of electrostatic interaction with other ions. On fig. 2 and 3 show examples of ionic bonds corresponding to the Kossel electron transfer model.
Rice. 2.
Rice. 3. Ionic bond in a molecule table salt(NaCl)
Here it is appropriate to recall some of the properties that explain the behavior of substances in nature, in particular, to consider the concept of acids And grounds.
Aqueous solutions All of these substances are electrolytes. They change color in different ways. indicators. The mechanism of action of indicators was discovered by F.V. Ostwald. He showed that the indicators are weak acids or bases, the color of which in the undissociated and dissociated states is different.
Bases can neutralize acids. Not all bases are soluble in water (for example, some organic compounds, not containing ‑ OH groups, in particular, triethylamine N (C 2 H 5) 3); soluble bases are called alkalis.
Aqueous solutions of acids enter into characteristic reactions:
a) with metal oxides - with the formation of salt and water;
b) with metals - with the formation of salt and hydrogen;
c) with carbonates - with the formation of salt, CO 2 and H 2 O.
The properties of acids and bases are described by several theories. In accordance with the theory of S.A. Arrhenius, an acid is a substance that dissociates to form ions H+ , while the base forms ions HE- . This theory does not take into account the existence of organic bases that do not have hydroxyl groups.
In line with proton Bronsted and Lowry's theory, an acid is a substance containing molecules or ions that donate protons ( donors protons), and the base is a substance consisting of molecules or ions that accept protons ( acceptors protons). Note that in aqueous solutions, hydrogen ions exist in a hydrated form, that is, in the form of hydronium ions H3O+ . This theory describes reactions not only with water and hydroxide ions, but also carried out in the absence of a solvent or with a non-aqueous solvent.
For example, in the reaction between ammonia NH 3 (weak base) and hydrogen chloride in the gas phase, solid ammonium chloride is formed, and in an equilibrium mixture of two substances there are always 4 particles, two of which are acids, and the other two are bases:
This equilibrium mixture consists of two conjugated pairs of acids and bases:
1)NH 4+ and NH 3
2) HCl And Cl ‑
Here, in each conjugated pair, the acid and base differ by one proton. Every acid has a conjugate base. strong acid A weak conjugate base corresponds to a weak acid, and a strong conjugate base corresponds to a weak acid.
The Bronsted-Lowry theory makes it possible to explain the unique role of water for the life of the biosphere. Water, depending on the substance interacting with it, can exhibit the properties of either an acid or a base. For example, in reactions with aqueous solutions of acetic acid, water is a base, and with aqueous solutions of ammonia, it is an acid.
1) CH 3 COOH + H 2 O ↔ H 3 O + + CH 3 SOO- . Here the acetic acid molecule donates a proton to the water molecule;
2) NH3 + H 2 O ↔ NH4 + + HE- . Here the ammonia molecule accepts a proton from the water molecule.
Thus, water can form two conjugated pairs:
1) H 2 O(acid) and HE- (conjugate base)
2) H 3 O+ (acid) and H 2 O(conjugate base).
In the first case, water donates a proton, and in the second, it accepts it.
Such a property is called amphiprotonity. Substances that can react as both acids and bases are called amphoteric. Such substances are often found in nature. For example, amino acids can form salts with both acids and bases. Therefore, peptides readily form coordination compounds with the metal ions present.
Thus, characteristic property ionic bond - the complete movement of a bunch of binding electrons to one of the nuclei. This means that there is a region between the ions where the electron density is almost zero.
The second type of connection iscovalent connection
Atoms can form stable electronic configurations by sharing electrons.
Such a bond is formed when a pair of electrons is shared one at a time. from each atom. In this case, the socialized bond electrons are distributed equally among the atoms. An example of a covalent bond is homonuclear diatomic H molecules 2 , N 2 , F 2. Allotropes have the same type of bond. O 2 and ozone O 3 and for a polyatomic molecule S 8 and also heteronuclear molecules hydrogen chloride Hcl, carbon dioxide CO 2, methane CH 4, ethanol WITH 2 H 5 HE, sulfur hexafluoride SF 6, acetylene WITH 2 H 2. All these molecules have the same common electrons, and their bonds are saturated and directed in the same way (Fig. 4).
For biologists, it is important that the covalent radii of atoms in double and triple bonds are reduced compared to a single bond.
Rice. 4. Covalent bond in the Cl 2 molecule.
Ionic and covalent types of bonds are two limiting cases of many existing types of chemical bonds, and in practice most of the bonds are intermediate.
Connections of two elements located at opposite ends of one or different periods systems of Mendeleev, predominantly form ionic bonds. As the elements approach each other within a period, the ionic nature of their compounds decreases, while the covalent character increases. For example, the halides and oxides of the elements on the left side periodic table form predominantly ionic bonds ( NaCl, AgBr, BaSO 4 , CaCO 3 , KNO 3 , CaO, NaOH), and the same compounds of the elements on the right side of the table are covalent ( H 2 O, CO 2, NH 3, NO 2, CH 4, phenol C6H5OH, glucose C 6 H 12 O 6, ethanol C 2 H 5 OH).
The covalent bond, in turn, has another modification.
In polyatomic ions and in complex biological molecules, both electrons can only come from one atom. It is called donor electron pair. An atom that socializes this pair of electrons with a donor is called acceptor electron pair. This type of covalent bond is called coordination (donor-acceptor, ordative) communication(Fig. 5). This type of bond is most important for biology and medicine, since the chemistry of the most important d-elements for metabolism is largely described by coordination bonds.
Pic. 5.
As a rule, in complex compound the metal atom acts as an electron pair acceptor; on the contrary, in ionic and covalent bonds, the metal atom is an electron donor.
The essence of the covalent bond and its variety - the coordination bond - can be clarified with the help of another theory of acids and bases, proposed by GN. Lewis. He somewhat expanded the semantic concept of the terms "acid" and "base" according to the Bronsted-Lowry theory. The Lewis theory explains the nature of the formation of complex ions and the participation of substances in nucleophilic substitution reactions, that is, in the formation of CS.
According to Lewis, an acid is a substance capable of forming a covalent bond by accepting an electron pair from a base. A Lewis base is a substance that has a lone pair of electrons, which, by donating electrons, forms a covalent bond with Lewis acid.
That is, the Lewis theory expands the range of acid-base reactions also to reactions in which protons do not participate at all. Moreover, the proton itself, according to this theory, is also an acid, since it is able to accept an electron pair.
Therefore, according to this theory, cations are Lewis acids and anions are Lewis bases. The following reactions are examples:
It was noted above that the subdivision of substances into ionic and covalent ones is relative, since there is no complete transition of an electron from metal atoms to acceptor atoms in covalent molecules. In compounds with an ionic bond, each ion is in the electric field of ions of the opposite sign, so they are mutually polarized, and their shells are deformed.
Polarizability determined by the electronic structure, charge and size of the ion; it is higher for anions than for cations. The highest polarizability among cations is for cations of larger charge and smaller size, for example, for Hg 2+ , Cd 2+ , Pb 2+ , Al 3+ , Tl 3+. Has a strong polarizing effect H+ . Since the effect of ion polarization is two-sided, it significantly changes the properties of the compounds they form.
The third type of connection -dipole-dipole connection
In addition to the listed types of communication, there are also dipole-dipole intermolecular interactions, also known as van der Waals .
The strength of these interactions depends on the nature of the molecules.
There are three types of interactions: permanent dipole - permanent dipole ( dipole-dipole attraction); permanent dipole - induced dipole ( induction attraction); instantaneous dipole - induced dipole ( dispersion attraction, or London forces; rice. 6).
Rice. 6.
Only molecules with polar covalent bonds have a dipole-dipole moment ( HCl, NH 3, SO 2, H 2 O, C 6 H 5 Cl), and the bond strength is 1-2 debye(1D \u003d 3.338 × 10 -30 coulomb meters - C × m).
In biochemistry, another type of bond is distinguished - hydrogen connection, which is a limiting case dipole-dipole attraction. This bond is formed by the attraction between a hydrogen atom and a small electronegative atom, most often oxygen, fluorine and nitrogen. With large atoms that have a similar electronegativity (for example, with chlorine and sulfur), the hydrogen bond is much weaker. The hydrogen atom is distinguished by one essential feature: when the binding electrons are pulled away, its nucleus - the proton - is exposed and ceases to be screened by electrons.
Therefore, the atom turns into a large dipole.
A hydrogen bond, unlike a van der Waals bond, is formed not only during intermolecular interactions, but also within one molecule - intramolecular hydrogen bond. Hydrogen bonds play an important role in biochemistry, for example, for stabilizing the structure of proteins in the form of an α-helix, or for the formation of a DNA double helix (Fig. 7).
Fig.7.
Hydrogen and van der Waals bonds are much weaker than ionic, covalent, and coordination bonds. The energy of intermolecular bonds is indicated in Table. 1.
Table 1. Energy of intermolecular forces
Note: The degree of intermolecular interactions reflect the enthalpy of melting and evaporation (boiling). Ionic compounds require much more energy to separate ions than to separate molecules. The melting enthalpies of ionic compounds are much higher than those of molecular compounds.
The fourth type of connection -metallic bond
Finally, there is another type of intermolecular bonds - metal: connection of positive ions of the lattice of metals with free electrons. This type of connection does not occur in biological objects.
From a brief review of the types of bonds, one detail emerges: an important parameter of an atom or ion of a metal - an electron donor, as well as an atom - an electron acceptor is its size.
Without going into details, we note that the covalent radii of atoms, the ionic radii of metals, and the van der Waals radii of interacting molecules increase as their atomic number in the groups of the periodic system increases. In this case, the values of the ion radii are the smallest, and the van der Waals radii are the largest. As a rule, when moving down the group, the radii of all elements increase, both covalent and van der Waals.
The most important for biologists and physicians are coordination(donor-acceptor) bonds considered by coordination chemistry.
Medical bioinorganics. G.K. Barashkov
First of all, consider the structure of the ammonia molecule NH 3 . As you already know, at the outer energy level, nitrogen atoms contain five electrons, of which three electrons are unpaired. It is they who are involved in the formation of three covalent bonds with three hydrogen atoms in the formation of an ammonia molecule NH 3 .
Three common electron pairs are shifted towards the more electronegative nitrogen atom, and since the ammonia molecule has the shape of a triangular pyramid (Fig. 128), a dipole arises as a result of the displacement of electron pairs, that is, a molecule with two poles.
Rice. 128.
The structure of the ammonia molecule
Ammonia molecules (in liquid ammonia) interact by binding to each other:
This special type of chemical intermolecular bond, as you already know, is called a hydrogen bond.
Ammonia is a colorless gas with a pungent odor, almost twice as light as air. Ammonia must not be inhaled long time because it is poisonous. This gas is easily liquefied at normal pressure and temperature -33.4 °C. When liquid ammonia evaporates from environment a lot of heat is absorbed, so ammonia is used in refrigeration.
Ammonia is highly soluble in water: at 20 ° C, about 710 volumes of ammonia dissolve in 1 volume of water (Fig. 129). Concentrated (25% by mass) aqueous ammonia is called aqueous ammonia or ammonia water, and a 10% ammonia solution used in medicine is known as ammonia. In an aqueous solution of ammonia, an unstable compound is formed - ammonia hydrate NH 3 H 2 O.
Rice. 129.
"Ammonia Fountain" (dissolving ammonia in water)
If a few drops of phenolphthalein are added to an ammonia solution, the solution will turn crimson, indicating alkaline environment. The alkaline reaction of aqueous solutions of ammonia is explained by the presence of hydroxide ions OH -:
If an ammonia solution dyed with phenolphthalein is heated, the color will disappear (why?).
Laboratory experiment No. 30
Studying the properties of ammonia
Ammonia reacts with acids to form ammonium salts. This interaction can be observed in the following experiment: a glass rod or a glass moistened with an ammonia solution is brought to another rod or glass moistened with hydrochloric acid, - thick white smoke will appear (Fig. 130):
Rice. 130.
"Smoke Without Fire"
So believe after this saying that there is no smoke without fire.
Both an aqueous solution of ammonia and ammonium salts contain a special ion - the ammonium cation NH + 4, which plays the role of a metal cation. The ammonium ion is formed as a result of the emergence of a covalent bond between the nitrogen atom, which has a free (lone) electron pair, and the hydrogen cation, which passes to ammonia from acid or water molecules:
In the formation of an ammonium ion, the donor of a free electron pair is the nitrogen atom in ammonia, and the acceptor is the hydrogen cation of an acid or water.
You can predict another chemical property of ammonia yourself if you pay attention to the degree of oxidation of nitrogen atoms in it, namely -3. Of course, ammonia is the strongest reducing agent, that is, its nitrogen atoms can only donate electrons, but not accept them. So, ammonia is able to oxidize or to free nitrogen (without the participation of a catalyst):
4NH 3 + 3O 2 \u003d 2N 2 + 6H 2 O,
or to nitric oxide (II) (in the presence of a catalyst):
In industry, ammonia is produced by synthesis from nitrogen and hydrogen (Fig. 131).
Rice. 131.
Industrial plant (a) and scheme for the industrial production of ammonia (b)
In the laboratory, ammonia is obtained by the action of slaked lime Ca (OH) 2 on ammonium salts, most often on ammonium chloride:
The gas is collected in a vessel turned upside down, and is recognized either by the smell, or by the blue of wet red litmus paper, or by the appearance of white smoke when a stick moistened with hydrochloric acid is introduced.
Ammonia and its salts are widely used in industry and technology, in agriculture, life. The main areas of their application are shown in Figure 132.
Rice. 132.
Application of ammonia and ammonium salts:
1,2 - in refrigeration units; 3 - production of mineral fertilizers; 4 - production nitric acid; 5 - for soldering; 6 - obtaining explosives; 7 - in medicine and in everyday life (ammonia)
New words and concepts
- The structure of the ammonia molecule.
- Hydrogen bond.
- Properties of ammonia: interaction with water, acids and oxygen.
- Donor-acceptor mechanism of ammonium ion formation.
- Obtaining, collecting and recognizing ammonia.
NH3 is one of the most famous and useful chemical substances. It has found wide application in the agricultural industry and not only. It has unique chemical properties due to which it is used in various industries.
What is NH3
NH 3 is known even to the most distant person from chemistry. It's ammonia. Ammonia (NH 3) is otherwise called hydrogen nitride and is at normal conditions A colorless gas with a distinct, characteristic odor. It is also worth noting that NH 3 gas (called ammonia) is almost twice as light as air!
In addition to gas, it can be a liquid at a temperature of about 70 ° C or exist as a solution (ammonia solution). A distinctive feature of liquid NH 3 is the ability to dissolve in itself the metals of the main subgroups of groups I and II of the table of elements of D. I. Mendeleev (that is, alkaline and alkaline earth metals), as well as magnesium, aluminium, europium and ytterbium. Unlike water, liquid ammonia does not interact with the above elements, but acts precisely as a solvent. This property allows metals to be isolated in their original form by evaporation of the solvent (NH 3). In the figure below, you can see what sodium dissolved in liquid ammonia looks like.
What does ammonia look like in terms of chemical bonds?
The scheme of ammonia (NH 3) and its spatial structure is most clearly shown by a triangular pyramid. The top of the "pyramid" of ammonia is the nitrogen atom (highlighted in blue), as seen in the image below.
The atoms in a substance called ammonia (NH 3 ) are linked by hydrogen bonds, just like in a water molecule. But it is very important to remember that the bonds in the ammonia molecule are weaker than in the water molecule. This explains why the melting and boiling points of NH 3 are lower when compared to H 2 O.
Chemical properties
The most common are 2 ways to obtain an NH 3 substance called ammonia. In industry, the so-called Haber process is used, the essence of which is the binding of atmospheric nitrogen and hydrogen (obtained from methane) by passing a mixture of these gases at high pressure over a heated catalyst.
In laboratories, the synthesis of ammonia is most often based on the interaction of concentrated ammonium chloride with solid sodium hydroxide.
Let's proceed to a direct consideration of the chemical properties of NH 3.
1) NH 3 acts as a weak base. That is why the following equation describing the interaction with water takes place:
NH 3 + H 2 O \u003d NH4 + + OH -
2) Also, its ability to react with acids and form the corresponding ammonium salts is based on the basic properties of NH 3:
NH3 + HNO 3 = NH 4 NO 3 (ammonium nitrate)
3) Earlier it was said that a certain group of metals dissolves in liquid ammonia. However, some metals are also able not only to dissolve, but to form compounds with NH 3 called amides:
Na (tv) + NH3 (g) = NaNH 2 + H 2
Na (tv) + NH3 (l) \u003d NaNH 2 + H 2 (the reaction is carried out in the presence of iron as a catalyst)
4) When NH 3 interacts with metals Fe 3+, Cr 3+, Al 3+, Sn 4+, Sn 2+, the corresponding metal hydroxides and ammonium cation are formed:
Fe 3+ + NH 3 + H 2 O \u003d Fe (OH) 3 + NH 4 +
5) The result of the interaction of NH 3 with metals Cu 2+, Ni 2+, Co 2+, Pd 2+, Pt 2+, Pt 4+ most often are the corresponding metal complexes:
Cu 2+ + NH 3 + H 2 O \u003d Cu (OH) 2 + NH 4 +
Cu (OH) 2 + NH 3 \u003d 2 + + OH -
Formation and further path of NH3 in the human body
It is well known that amino acids are an integral part of biochemical processes in the human body. They are the main source of NH 3, a substance called ammonia, the result of their oxidative deamination (most often). Unfortunately, ammonia is toxic for the human body; the above-mentioned ammonium cation (NH 4 +) is easily formed from it and accumulates in cells. Subsequently, there is a slowdown in the most important biochemical cycles, and as a result, a drop in the level of ATP produced.
It is easy to guess that the body needs mechanisms to bind and neutralize the released NH 3 . The diagram below shows the sources and some of the binding products of ammonia in the human body.
So, in short, the neutralization of ammonia occurs through the formation of its transport forms in tissues (for example, glutamine and alanine), by excretion in the urine, with the help of urea biosynthesis, which is the main natural way of neutralizing NH 3 in the human body.
The use of NH3 - a substance called ammonia
In modern times, liquid ammonia is the most concentrated and cheapest nitrogen fertilizer, which is used in agriculture for ammonization of coarse soils and peat. When liquid ammonia is applied to the soil, an increase in the number of microorganisms occurs, but no negative consequences as, for example, from solid fertilizers. The figure below shows one of the possible installations for liquefying gaseous ammonia with liquid nitrogen.
Evaporating, liquid ammonia absorbs a lot of heat from the environment, causing cooling. This property is used in refrigeration plants to obtain artificial ice when storing perishable foodstuffs. In addition, it is used to freeze the soil during the construction of underground structures. Aqueous solutions of ammonia are used in the chemical industry (it is an industrial non-aqueous solvent), laboratory practice (for example, as a solvent in the electrochemical production of chemical products), medicine and household use.
