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5.1. Covalent bond
A chemical bond is formed when two or more atoms come together if, as a result of their interaction, the total energy of the system decreases. The most stable electronic configurations of the outer electron shells of atoms are those 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 unfinished. 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.
Covalent is a chemical bond between two atoms, which is formed by electron pairs that simultaneously belong to these two atoms.
There are two mechanisms for the formation of covalent bonds: exchange and donor-acceptor.
5.1.1. Exchange mechanism of covalent bond formation
Exchange mechanism The formation of a covalent bond is realized due to the overlap 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, a 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 for the formation of covalent bonds
With the donor-acceptor mechanism of covalent bond formation, the bond is also formed using 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 ammonium ion NH 4 + during the interaction of ammonia NH 3 with the 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 lone. This electron pair can form a bond with a hydrogen ion that has an unoccupied orbital. The result is ammonium ion NH 4 +, Fig. 5.2.
Rice. 5.2. The appearance of a donor-acceptor bond during the formation of ammonium ion
It should be noted that the four covalent N–H bonds existing in the NH 4 + ion are equivalent. In the ammonium ion it is impossible to identify a 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 nonpolar. Examples of molecules with a non-polar covalent bond are H2, Cl2, O2, N2, etc.
In the case of a polar covalent bond, the shared electron pair is shifted to the atom with higher electronegativity. This type of bond is realized in molecules formed by different atoms. A polar covalent bond occurs in molecules of HCl, HBr, CO, NO, etc. For example, the formation of a polar covalent bond in a HCl molecule can be represented by a diagram, 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 molecular structure
A measure of the polarity of a bond is its dipole moment μ:
μ = e l,
Where e– electron charge, l– the distance between the centers of positive and negative charges.
Dipole moment is a vector quantity. The concepts of “bond dipole moment” and “molecule dipole moment” 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 nonpolar, since the dipole moments of the bonds cancel each other out (Fig. 5.4). The dipole moment of the carbon dioxide molecule is m = 0.
In the angular H2O 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 covalent bonds
The valence 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 in the outer layer, therefore in compounds such as O 2, H 2 O, H 2 S and SCl 2 they are divalent.
Since, in addition to ordinary 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. A quantitative measure of valency is the number of chemical bonds through 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 secondary subgroup of the first group Cu, Ag, Au, whose valence in compounds is greater than one. The valence electrons primarily include the electrons of the outer layers, however, for elements of side subgroups, the electrons of the penultimate (pre-outer) layers also take part in the formation of a chemical bond.
5.1.6. Valence of elements in normal and excited states
The valency of most chemical elements depends on whether these elements are in a 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 the 1s electron to the 2p orbital 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 trivalent lithium compounds.
Configuration of the outer electronic layer of elements of the beryllium subgroup ns 2. This means that in the outer electron layer of these elements in the ns cell orbital there are two electrons with opposite spins. Elements of the beryllium subgroup do not contain unpaired electrons, so their valence 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
Let's consider the electronic configuration of the boron atom in the ground state: 1s 2 2s 2 2p 1. The boron atom in the ground state contains one unpaired electron (Fig. 5.5), i.e. it is monovalent. However, boron is not characterized by the formation of compounds in which it is monovalent. When a boron atom is excited, one 2s electron transitions to a 2p orbital (Fig. 5.5). A 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 normal and excited states
The energy expended on the transition of an atom to an excited state within one energy level, as a rule, is more than compensated 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, resulting in the formation of the – ion, 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
Let's consider the electronic structure of the nitrogen atom (Fig. 5.7).
Rice. 5.7. Distribution of electrons in the orbitals of the nitrogen atom
From the presented diagram it is clear 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 a lone electron pair of outer electrons 2s 2 to an atom having a free orbital (acceptor). As a result, a fourth chemical bond of the nitrogen atom appears, as is the case, for example, in the ammonium ion (Fig. 5.2). Thus, the maximum covalency (the number of covalent bonds formed) of a 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 the atoms of nitrogen, oxygen and fluorine, the atoms of phosphorus, sulfur and chlorine located 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 orbitals for a phosphorus atom in an excited state
In the excited state, sulfur, in addition to a valence of two, also exhibits a valence of four and six. In this case, 3p and 3s electrons are sequentially paired (Fig. 5.9).
Rice. 5.9. Valence possibilities of a sulfur atom in an excited state
In the excited state, for all elements of the main subgroup of group V, except fluorine, sequential pairing 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
Covalent bonds typically form between nonmetal atoms. The main characteristics of a covalent bond are length, energy and direction.
Covalent bond length
The length of a bond is the distance between the nuclei of the atoms forming 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 molecules A 2 and B 2:
.
From top to bottom along the subgroups of the periodic system of elements, the length of the chemical bond increases, since the radii of the atoms increase in this direction (Table 5.1). As the bond multiplicity increases, its length decreases.
Table 5.1.
Length of some chemical bonds
Chemical bond |
Link length, pm |
Chemical bond |
Link length, pm |
C – C |
|||
Communication energy
A measure of bond strength is the bond energy. Communication energy determined by the energy required to break a bond and remove the atoms forming that bond to an infinitely large distance from each other. The covalent bond is very strong. Its energy ranges from several tens to several hundred kJ/mol. For an IСl 3 molecule, for example, the Ebond is ≈40, and for N 2 and CO molecules the Ebond is ≈1000 kJ/mol.
From top to bottom along 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). As the bond multiplicity increases, its energy increases (Table 5.2).
Table 5.2.
Energies of some chemical bonds
Chemical bond |
Communication energy, |
Chemical bond |
Communication energy, |
C – C |
|||
Saturation and directionality of covalent bonds
The most important properties of a covalent bond are its saturation and directionality. Saturability can be defined as the ability of atoms to form a limited number of covalent bonds. Thus, 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 direction 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, which occupy two orbitals located at right angles to each other. Hydrogen atoms have unpaired 1s electrons. The angle between the bonds formed by p-electrons should be close to the angle between the orbitals of 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, whose orbitals are located in three mutually perpendicular directions, participate in the formation of the NH 3 molecule. Therefore, the three N–H bonds should be located 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 between the angles between the bonds and the theoretical values is due, as in the case of the water molecule, to the mutual repulsion of hydrogen atoms. In addition, the presented schemes do not take into account the possibility of the participation of two electrons in the 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
Let's consider the formation of the BeC1 2 molecule. A beryllium atom 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 the BCl 3 and CCl 4 molecules - all bonds in these molecules are equivalent. The BC1 3 molecule has a flat structure, CC1 4 has a tetrahedral structure.
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 replacing several atomic orbitals, which do not differ very much in their energy, with the same number of equivalent orbitals, called hybrid ones. These hybrid orbitals are composed of atomic orbitals as a result of their linear combination.
According to L. Pauling, when chemical bonds are formed by an atom having electrons of different 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 different types, in which their alignment in shape and energy occurs . As a result, hybrid orbitals are formed that have an asymmetric shape and are highly elongated on one side of the nucleus. It is important to emphasize that the hybridization model is used when electrons of different types, for example s and p, are involved in the formation of bonds.
5.1.8.2. Various types of atomic orbital hybridization
sp hybridization
Hybridization of one s- and one R- orbitals ( sp- hybridization) is realized, for example, during the formation of beryllium chloride. As shown above, in an excited state, a Be atom has two unpaired electrons, one of which occupies the 2s orbital, and the other occupies 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 minimal 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. Diagram of electron orbital overlap during the formation of a BeCl 2 molecule
The structure of the acetylene molecule; sigma and pi bonds
Let's consider a diagram of the overlap of electronic orbitals during the formation of an acetylene molecule. In an acetylene molecule, each carbon atom is in an 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 formation of s-bonds in an acetylene molecule
A σ bond is a bond formed as a result of overlapping electron orbitals along a 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 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 the line connecting the nuclei of atoms.
Rice. 5.14. Scheme of the formation of σ - and π - bonds in the acetylene molecule.
Thus, in the acetylene molecule, a triple bond is formed between the 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. A boron atom in an excited state on the outer electron layer contains one s-electron and two p-electrons, 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 flat shape.
Rice. 5.15. Flat structure of the BCl 3 molecule
sp 3 - hybridization
The valence orbitals of the carbon atom (s, р x, р y, р 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 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 that, in addition to bonding ones, also contain lone pairs of electrons. In water and ammonia molecules, the total number of electron pairs of the central atom (O and N) is four. At the same time, a water molecule has two, and an ammonia molecule has one lone pair of electrons. The formation of chemical bonds in these molecules can be explained by assuming that lone pairs of electrons can also fill hybrid orbitals. Lone electron pairs take up much more space in space than bonding ones. 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. Establishing the type of hybridization and determining 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 lone pairs of electrons, is determined by the number of sigma bonds. If there are two such bonds, sp-hybridization occurs, three - sp 2 -hybridization, four - sp 3 -hybridization. Lone 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. in elements of the main subgroups II - IV groups.
2. If the central atom contains lone 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 lone electron pairs. Hybridization involving lone electron pairs occurs in molecules formed by atoms of nitrogen, phosphorus, oxygen, sulfur, 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.
Bond angles, geometric shape of molecules depending on the number of hybrid orbitals and the type of hybridization of the central atom
5.2. Ionic bond
Ionic bonding occurs through 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 alkali metal and halogen atoms.
Let us consider the occurrence 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 clear that to complete the outer level, it is easier for the sodium atom to give up one electron than to add seven, and it is easier for the chlorine atom to add one, than to give away seven. In chemical reactions, the sodium atom gives up one electron, and the chlorine atom takes it. As a result, the electronic shells of sodium and chlorine atoms are transformed into stable electronic 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 – - 1s 2 2s 2 2p 6 3s 2 3p 6). The electrostatic interaction of ions leads to the formation of a NaCl molecule.
Basic characteristics of ionic bonds and properties of ionic compounds
1. An ionic bond is a strong chemical bond. The energy of this bond is on the order of 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 opposite sign does not lead to complete mutual compensation of their force fields.
4. During the formation of molecules with an ionic bond, complete transfer of electrons does not occur, therefore, one hundred percent ionic bonds do not exist in nature. In the NaCl molecule, the chemical bond is only 80% ionic.
5. Compounds with ionic bonds are crystalline solids that have high melting and boiling points.
6. Most ionic compounds are soluble in water. Solutions and melts of ionic compounds conduct electric current.
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, metal cations are located in the nodes of their crystal lattice, and 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 causes each cation to interact with this electron. Thus, metallic bonding is the bonding between positive ions in metal crystals, which occurs through the attraction of electrons moving freely throughout the crystal.
Since the valence electrons in a metal are evenly distributed throughout the crystal, a metallic bond, like an ionic bond, is a non-directional bond. Unlike a covalent bond, a metallic bond is an unsaturated bond. From covalent bond metal connection It also differs in durability. The energy of a metallic bond is approximately 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 the 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 form 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 nature. In this case, the electron pair donor is an atom of a strongly electronegative element (F, O, N), and the acceptor is the hydrogen atoms connected to these atoms. Like covalent bonds, hydrogen bonds are characterized by focus in space and saturability.
Hydrogen bonds are usually denoted by dots: H ··· F. The stronger the hydrogen bond, the greater the electronegativity of the partner atom and the smaller its size. It is characteristic primarily of fluorine compounds, as well as oxygen, to a lesser extent nitrogen, and to an even lesser extent chlorine and sulfur. The energy of the hydrogen bond also changes accordingly (Table 5.4).
Table 5.4.
Average values of hydrogen bond energies
Intermolecular and intramolecular hydrogen bonding
Thanks to hydrogen bonds, molecules combine into dimers and more complex associates. For example, the formation of a formic acid dimer can be represented by the following diagram (Fig. 5.18).
Rice. 5.18. Formation of intermolecular hydrogen bonds in formic acid
Long chains of (H 2 O) n associates 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 H2O molecule can form four hydrogen bonds, but an HF molecule can form only two.
Hydrogen bonds can occur both between different molecules (intermolecular hydrogen bonding) and within a molecule (intramolecular hydrogen bonding). Examples of the formation of intramolecular bonds for some organic substances are presented in Fig. 5.20.
Rice. 5.20. Formation of intramolecular hydrogen bonds in molecules of various organic compounds
The influence of hydrogen bonding on the properties of substances
The most convenient indicator of the existence of intermolecular hydrogen bonds is the boiling point of a substance. The higher boiling point of water (100 o C compared to hydrogen compounds of elements of the oxygen subgroup (H 2 S, H 2 Se, H 2 Te) is explained by the presence of hydrogen bonds: additional energy must be expended to destroy intermolecular hydrogen bonds in water.
Hydrogen bonding 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 causes the deoxyribonucleic acid (DNA) molecule to be folded into a double helix in water.
Hydrogen bonding also plays an important role in dissolution processes, since solubility also depends on the ability of a compound to form hydrogen bonds with the solvent. As a result, substances containing OH groups such as sugar, glucose, alcohols, and carboxylic acids are, as a rule, highly soluble in water.
5.5. Types of crystal lattices
Solids usually 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 a given lattice. The points at which particles are located are called crystal 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 connected to each other by covalent bonds. Substances that have an atomic lattice include diamond, silicon, carbides, silicides, etc. In the structure of an atomic crystal it is impossible to isolate individual molecules; the entire crystal is considered as one giant molecule. The structure of diamond is shown in Fig. 5.21. Diamond is made up of carbon atoms, each of which is bonded to four neighboring atoms. Due to the fact that covalent bonds are strong, all substances with atomic lattices are refractory, hard and low-volatile. They are slightly soluble in water.
Rice. 5.21. Diamond crystal lattice
5.5.2. Molecular crystal lattice
At the nodes of molecular crystal lattices there are molecules connected to each other by weak intermolecular forces. Therefore, substances with a molecular lattice have low hardness, they are fusible, 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 containing ions at their nodes are called ionic. They are formed by substances with ionic bonds, 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, therefore substances with an ionic lattice have low volatility and 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, the light balls are Na + ions, the dark balls are Cl – ions. On the left in Fig. Figure 5.22 shows the unit cell of NaCl.
5.5.4. Metal crystal lattice
Metals in the solid state form metallic crystal lattices. The sites of such lattices contain positive metal ions, and valence electrons move freely between them. The electrons electrostatically attract cations, thereby imparting stability to the metal lattice. This lattice structure determines the high thermal conductivity, electrical conductivity and plasticity of metals - during mechanical deformation there is no breaking of bonds and destruction of the crystal, since the ions that make it up seem to float in a cloud of electron gas. In Fig. Figure 5.23 shows the sodium crystal lattice.
Rice. 5.23. Sodium crystal lattice
DEFINITION
Ammonia- hydrogen nitride.
Formula – NH 3. Molar mass – 17 g/mol.
Physical properties of ammonia
Ammonia (NH 3) is a colorless gas with a pungent odor (the smell of “ammonia”), lighter than air, highly soluble in water (one volume of water will dissolve up to 700 volumes of ammonia). The concentrated ammonia solution contains 25% (mass) ammonia and has a density of 0.91 g/cm 3 .
The bonds between atoms in the ammonia molecule are covalent. General view of the AB 3 molecule. All valence orbitals of the nitrogen atom enter into hybridization, therefore, the type of hybridization of the ammonia molecule is sp 3. Ammonia has a geometric structure of the AB 3 E type - a trigonal pyramid (Fig. 1).
Rice. 1. The structure of the ammonia molecule.
Chemical properties of ammonia
Chemically, ammonia is quite active: it reacts with many substances. The oxidation degree of nitrogen in ammonia “-3” is minimal, so ammonia exhibits only reducing properties.
When ammonia is heated with halogens, heavy metal oxides and oxygen, nitrogen is formed:
2NH 3 + 3Br 2 = N 2 + 6HBr
2NH 3 + 3CuO = 3Cu + N 2 + 3H 2 O
4NH 3 +3O 2 = 2N 2 + 6H 2 O
In the presence of a catalyst, ammonia can be oxidized to nitrogen oxide (II):
4NH 3 + 5O 2 = 4NO + 6H 2 O (catalyst - platinum)
Unlike hydrogen compounds of non-metals of groups VI and VII, ammonia does not exhibit acidic properties. However, hydrogen atoms in its molecule are still capable of being replaced by metal atoms. When hydrogen is completely replaced by a metal, compounds called nitrides are formed, which can also be obtained by direct interaction of nitrogen with the metal at high temperatures.
The main properties of ammonia are due to the presence of a lone pair of electrons on the nitrogen atom. A solution of ammonia in water is alkaline:
NH 3 + H 2 O ↔ NH 4 OH ↔ NH 4 + + OH —
When ammonia interacts with acids, ammonium salts are formed, which decompose when heated:
NH 3 + HCl = NH 4 Cl
NH 4 Cl = NH 3 + HCl (when heated)
Ammonia production
There are industrial and laboratory methods for producing ammonia. In the laboratory, ammonia is obtained by the action of alkalis on solutions of ammonium salts when heated:
NH 4 Cl + KOH = NH 3 + KCl + H 2 O
NH 4 + + OH - = NH 3 + H 2 O
This reaction is qualitative for ammonium ions.
Application of ammonia
Ammonia production is one of the most important technological processes worldwide. About 100 million tons of ammonia are produced annually in the world. Ammonia is released in liquid form or in the form of a 25% aqueous solution - ammonia water. The main areas of use of ammonia are the production of nitric acid (subsequent production of nitrogen-containing mineral fertilizers), ammonium salts, urea, hexamine, synthetic fibers (nylon and nylon). Ammonia is used as a refrigerant in industrial refrigeration units and as a bleaching agent in the cleaning and dyeing of cotton, wool and silk.
Examples of problem solving
EXAMPLE 1
| Exercise | What is the mass and volume of ammonia that will be required to produce 5 tons of ammonium nitrate? |
| Solution | Let us write the equation for the reaction of producing ammonium nitrate from ammonia and nitric acid: NH 3 + HNO 3 = NH 4 NO 3 According to the reaction equation, the amount of ammonium nitrate substance is equal to 1 mol - v(NH 4 NO 3) = 1 mol. Then, the mass of ammonium nitrate calculated from the reaction equation: m(NH 4 NO 3) = v(NH 4 NO 3) × M(NH 4 NO 3); m(NH 4 NO 3) = 1×80 = 80 t According to the reaction equation, the amount of ammonia substance is also equal to 1 mol - v(NH 3) = 1 mol. Then, the mass of ammonia calculated by the equation: m(NH 3) = v(NH 3)×M(NH 3); m(NH 3) = 1×17 = 17 t Let's make a proportion and find the mass of ammonia (practical): x g NH 3 – 5 t NH 4 NO 3 17 t NH 3 – 80 t NH 4 NO 3 x = 17×5/80 = 1.06 m(NH 3) = 1.06 t Let’s make a similar proportion to find the volume of ammonia: 1.06 g NH 3 – x l NH 3 17 t NH 3 – 22.4×10 3 m 3 NH 3 x = 22.4×10 3 ×1.06 /17 = 1.4×10 3 V(NH 3) = 1.4 × 10 3 m 3 |
| Answer | Ammonia mass - 1.06 t, ammonia volume - 1.4×10 m |
First of all, let's 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 participate in the formation of three covalent bonds with three hydrogen atoms during the formation of the 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), as a result of the displacement of electron pairs, a dipole appears, i.e. a molecule with two poles.
Rice. 128.
The structure of the ammonia molecule
Ammonia molecules (in liquid ammonia) interact by bonding with 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 should not be inhaled for long periods of time as it is poisonous. This gas easily liquefies at normal pressure and a temperature of -33.4 °C. When liquid ammonia evaporates from the environment, a lot of heat is absorbed, which is why ammonia is used in refrigeration units.
Ammonia is highly soluble in water: at 20 °C, about 710 volumes of ammonia dissolve in 1 volume of water (Fig. 129). A concentrated (25% by weight) aqueous solution of 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, a weak compound is formed - ammonia hydrate NH 3 H 2 O.
Rice. 129.
“Ammonia fountain” (dissolving ammonia in water)
If you add a few drops of phenolphthalein to an ammonia solution, the solution will turn crimson, indicating an alkaline environment. The alkaline reaction of aqueous solutions of ammonia is explained by the presence of hydroxide ions OH -:
If an ammonia solution colored 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: bring a glass rod or glass moistened with an ammonia solution 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 - ammonium cation NH + 4, which plays the role of a metal cation. The ammonium ion is formed as a result of the formation of a covalent bond between a nitrogen atom having a free (lone) electron pair and a hydrogen cation, which passes to ammonia from acid or water molecules:
When an ammonium ion is formed, 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 oxidation state of nitrogen atoms in it, namely -3. Of course, ammonia is the strongest reducing agent, that is, its nitrogen atoms can only give up electrons, but not accept them. Thus, ammonia can be oxidized either to free nitrogen (without the participation of a catalyst):
4NH 3 + 3O 2 = 2N 2 + 6H 2 O,
or to nitrogen oxide (II) (in the presence of a catalyst):
In industry, ammonia is produced by synthesis from nitrogen and hydrogen (Fig. 131).
Rice. 131.
Industrial installation (a) and scheme for 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 ammonium chloride:
The gas is collected in a vessel turned upside down, and is recognized either by smell, or by the blueness 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, agriculture, and everyday life. Their main areas of 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 of nitric acid; 5 - for soldering; 6 - production of 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 for the formation of ammonium ion.
- Receiving, collecting and recognizing ammonia.
7.11. The structure of substances with covalent bonds
Substances in which, of all types of chemical bonds, only a covalent one is present, are divided into two unequal groups: molecular (very many) and non-molecular (much less).
Crystals of solid molecular substances consist of molecules weakly bound together by the forces of intermolecular interaction of molecules. Such crystals do not have high strength and hardness (think ice or sugar). Their melting and boiling points are also low (see Table 22).
Table 22. Melting and boiling points of some molecular substances
Substance |
Substance |
||||
| H 2 | – 259 | – 253 | BR 2 | – 7 | 58 |
| N 2 | – 210 | – 196 | H2O | 0 | 100 |
| HCl | – 112 | – 85 | P 4 | 44 | 257 |
| NH 3 | – 78 | – 33 | C 10 H 8 (naphthalene) | 80 | 218 |
| SO 2 | – 75 | – 10 | S 8 | 119 |
Unlike their molecular counterparts, non-molecular substances with covalent bonds form very hard crystals. Diamond crystals (the hardest substance) belong to this type.
In a diamond crystal (Fig. 7.5), each carbon atom is connected to four other carbon atoms by simple covalent bonds (sp 3 hybridization). The carbon atoms form a three-dimensional framework. Essentially the entire diamond crystal is one huge and very strong molecule.
Silicon crystals, widely used in radio electronics and electronic engineering, have the same structure.
If you replace half of the carbon atoms in diamond with silicon atoms without disturbing the framework structure of the crystal, you will get a crystal of silicon carbide SiC - also a very hard substance used as an abrasive material. Ordinary quartz sand (silicon dioxide) also belongs to this type of crystalline substance. Quartz is a very hard substance; Under the name "emery" it is also used as an abrasive material. The quartz structure is easily obtained by inserting oxygen atoms between every two silicon atoms in a silicon crystal. In this case, each silicon atom will be associated with four oxygen atoms, and each oxygen atom with two silicon atoms.
Crystals of diamond, silicon, quartz and similar structures are called atomic crystals.
An atomic crystal is a crystal consisting of atoms of one or more elements linked by chemical bonds.
A chemical bond in an atomic crystal can be covalent or metallic.
As you already know, any atomic crystal, like an ionic crystal, is a huge “supermolecule”. The structural formula of such a “supermolecule” cannot be written down - you can only show its fragment, for example:
Unlike molecular substances, substances that form atomic crystals are among the most refractory (see table 23.).
Table 23. Melting and boiling points of some non-molecular substances With covalent bonds
Such high melting temperatures are quite understandable if we remember that when these substances melt, it is not weak intermolecular bonds that are broken, but strong chemical bonds. For the same reason, many substances that form atomic crystals do not melt when heated, but decompose or immediately transform into a vapor state (sublimate), for example, graphite sublimes at 3700 o C.
| Silicon – Si. Very hard, brittle silicon crystals look like metal, but it is nevertheless a non-metal. Based on the type of electrical conductivity, this substance is classified as a semiconductor, which determines its enormous importance in the modern world. Silicon is the most important semiconductor material. Radios, televisions, computers, modern telephones, electronic watches, solar panels and many other household and industrial devices contain transistors, microcircuits and photocells made from single crystals of high-purity silicon as the most important structural elements. Technical silicon is used in steel production and non-ferrous metallurgy. In terms of its chemical properties, silicon is a fairly inert substance; it reacts only at high temperatures. Silicon dioxide – SiO 2 . Another name for this substance is silica. Silicon dioxide occurs in nature in two forms: crystalline and amorphous. Many semi-precious and ornamental stones are varieties of crystalline silicon dioxide (quartz): rock crystal, jasper, chalcedony, agate. and opal is an amorphous form of silica. Quartz is very widespread in nature, because dunes in deserts and sandbanks of rivers and seas are all quartz sand. Quartz is a colorless crystalline, very hard and refractory substance. It is inferior in hardness to diamond and corundum, but, nevertheless, is widely used as an abrasive material. Quartz sand is widely used in construction and the building materials industry. Quartz glass is used to make laboratory glassware and scientific instruments because it does not crack under sudden temperature changes. In terms of its chemical properties, silicon dioxide is an acidic oxide, but it reacts with alkalis only when fused. At high temperatures, silicon dioxide and graphite are used to produce silicon carbide - carborundum. Carborundum is the second hardest substance after diamond; it is also used to make grinding wheels and “sandpaper”. |
7.12. Polarity of a covalent bond. Electronegativity
Recall that isolated atoms of different elements have different propensities to both give up and accept electrons. These differences persist after the formation of a covalent bond. That is, atoms of some elements tend to attract the electron pair of a covalent bond to themselves more strongly than atoms of other elements.
Consider a molecule HCl.
Using this example, let's see how we can estimate the displacement of the electron communication cloud using molar ionization energies and means to the electron. 1312 kJ/mol, and 1251 kJ/mol - the difference is insignificant, approximately 5%. 73 kJ/mol, and 349 kJ/mol - here the difference is much greater: the electron affinity energy of the chlorine atom is almost five times greater than that for the hydrogen atom. From this we can conclude that the electron pair of the covalent bond in the hydrogen chloride molecule is largely shifted towards the chlorine atom. In other words, the bonding electrons spend more time near the chlorine atom than near the hydrogen atom. This uneven distribution of electron density leads to a redistribution of electrical charges inside the molecule. Partial (excess) charges arise on the atoms; on the hydrogen atom it is positive, and on the chlorine atom it is negative.
In this case, the bond is said to be polarized, and the bond itself is called a polar covalent bond.
If the electron pair of a covalent bond is not displaced to any of the bonded atoms, that is, the bond electrons equally belong to the bonded atoms, then such a bond is called a nonpolar covalent bond.
The concept of "formal charge" in the case of a covalent bond is also applicable. Only in the definition we should not be talking about ions, but about atoms. In general, the following definition can be given.
In molecules in which covalent bonds are formed only by an exchange mechanism, the formal charges of the atoms are equal to zero. Thus, in the HCl molecule, the formal charges on both the chlorine and hydrogen atoms are zero. Consequently, in this molecule the real (effective) charges on the chlorine and hydrogen atoms are equal to the partial (excess) charges.
It is not always easy to determine the sign of the partial charge on an atom of one or another element in a molecule based on the molar ionization energies and affinity for the electrode, that is, to estimate in which direction the electron pairs of bonds are shifted. Usually, for these purposes, another energy characteristic of an atom is used - electronegativity.
Currently, there is no single, generally accepted designation for electronegativity. It can be denoted by the letters E/O. There is also no single, generally accepted method for calculating electronegativity. In a simplified way, it can be represented as half the sum of the molar ionization energies and electron affinity - this was one of the first ways to calculate it.
The absolute values of electronegativity of atoms of various elements are used very rarely. The most commonly used is relative electronegativity, denoted by c. Initially, this value was defined as the ratio of the electronegativity of an atom of a given element to the electronegativity of a lithium atom. Subsequently, the methods of its calculation changed somewhat.
Relative electronegativity is a dimensionless quantity. Its values are given in Appendix 10.
Since relative electronegativity depends primarily on the ionization energy of the atom (electron affinity energy is always much lower), then in a system of chemical elements it changes approximately the same as the ionization energy, that is, it increases diagonally from cesium (0.86) to fluorine (4.10). The values of the relative electronegativity of helium and neon given in the table have no practical significance, since these elements do not form compounds.
Using the electronegativity table, you can easily determine towards which of the two atoms the electrons connecting these atoms are shifted, and, therefore, the signs of the partial charges arising on these atoms.
| H2O | The connection is polar | |||
| H 2 | Atoms are the same | H--H | The connection is non-polar | |
| CO2 | The connection is polar | |||
| Cl2 | Atoms are the same | Cl--Cl | The connection is non-polar | |
| H2S | The connection is polar |
Thus, in the case of the formation of a covalent bond between atoms of different elements, such a bond will always be polar, and in the case of the formation of a covalent bond between atoms of the same element (in simple substances), the bond is in most cases non-polar.
The greater the difference in electronegativity of the bonded atoms, the more polar the covalent bond between these atoms turns out to be.
| Hydrogen sulfide H 2 S– a colorless gas with a characteristic odor characteristic of rotten eggs; poisonous. It is thermally unstable and decomposes when heated. Hydrogen sulfide is slightly soluble in water; its aqueous solution is called hydrosulfide acid. Hydrogen sulfide provokes (catalyzes) corrosion of metals; it is this gas that is “to blame” for the darkening of silver. It is naturally found in some mineral waters. In the process of life, it is formed by some bacteria. Hydrogen sulfide is destructive to all living things. A hydrogen sulfide layer was discovered in the depths of the Black Sea and causes concern to scientists: the life of marine inhabitants there is under constant threat. |
POLAR COVALENT BOND, NON-POLAR COVALENT BOND, ABSOLUTE ELECTRONEGATIVITY, RELATIVE ELECTRONEGATIVITY.
1. Experiments and subsequent calculations showed that the effective charge of silicon in silicon tetrafluoride is +1.64 e, and of xenon in xenon hexafluoride +2.3 e. Determine the values of the partial charges on the fluorine atoms in these compounds. 2. Make up the structural formulas of the following substances and, using the notations " " and " ", characterize the polarity of covalent bonds in the molecules of these compounds: a) CH 4, CCl 4, SiCl 4; b) H 2 O, H 2 S, H 2 Se, H 2 Te; c) NH 3, NF 3, NCl 3; d) SO 2, Cl 2 O, OF 2.
3.Using the electronegativity table, indicate in which of the compounds the bond is more polar: a) CCl 4 or SiCl 4 ; b) H 2 S or H 2 O; c) NF 3 or NCl 3; d) Cl 2 O or OF 2.
7.13. Donor-acceptor mechanism of bond formation
In the previous paragraphs, you learned in detail about two types of bonds: ionic and covalent. Recall that an ionic bond is formed when an electron is completely transferred from one atom to another. Covalent - when sharing unpaired electrons of bonded atoms.
In addition, there is another mechanism for bond formation. Let's consider it using the example of the interaction of an ammonia molecule with a boron trifluoride molecule:
As a result, both covalent and ionic bonds arise between the nitrogen and boron atoms. In this case, the nitrogen atom is donor electron pair ("gives" it for the formation of a bond), and the boron atom - acceptor(“accepts” it when forming a connection). Hence the name of the mechanism for the formation of such a connection - “ donor-acceptor".
When a bond is formed using the donor-acceptor mechanism, both a covalent bond and an ionic bond are formed simultaneously.
Of course, after the formation of a bond, due to the difference in the electronegativity of the bonded atoms, polarization of the bond occurs and partial charges arise, reducing the effective (real) charges of the atoms.
Let's look at other examples.
If there is a highly polar hydrogen chloride molecule next to the ammonia molecule, in which there is a significant partial charge on the hydrogen atom, then in this case the role of the electron pair acceptor will be played by the hydrogen atom. Its 1 s-AO, although not completely empty, like the boron atom in the previous example, the electron density in the cloud of this orbital is significantly reduced.
The spatial structure of the resulting cation is ammonium ion NH 4 is similar to the structure of the methane molecule, that is, all four N-H bonds are exactly the same.
The formation of ionic crystals of ammonium chloride NH 4 Cl can be observed by mixing ammonia gas with hydrogen chloride gas:
NH 3 (g) + HCl (g) = NH 4 Cl (cr)
Not only the nitrogen atom can be an electron pair donor. It could be, for example, the oxygen atom of a water molecule. A water molecule will interact with the same hydrogen chloride as follows:
The resulting H3O cation is called oxonium ion and, as you will soon learn, is of great importance in chemistry.
In conclusion, let us consider the electronic structure of the carbon monoxide (carbon monoxide) CO molecule:
In addition to three covalent bonds (triple bond), it also contains an ionic bond.
Conditions for bond formation according to the donor-acceptor mechanism:
1) the presence of a lone pair of valence electrons in one of the atoms;
2) the presence of a free orbital on the valence sublevel of another atom.
The donor-acceptor mechanism of bond formation is quite widespread. It occurs especially often during the formation of compounds d-elements. Almost everyone's atoms d-elements have many empty valence orbitals. Therefore, they are active acceptors of electron pairs.
DONOR-ACCEPTOR MECHANISM OF BOND FORMATION, AMMONIUM ION, OXONIUM ION, CONDITIONS FOR BOND FORMATION BY DONOR-ACCEPTOR MECHANISM.
1.Make reaction equations and formation schemes
a) ammonium bromide NH 4 Br from ammonia and hydrogen bromide;
b) ammonium sulfate (NH 4) 2 SO 4 from ammonia and sulfuric acid.
2. Create reaction equations and interaction schemes for a) water with hydrogen bromide; b) water with sulfuric acid.
3.Which atoms in the four previous reactions are donors of an electron pair, and which are acceptors? Why? Explain your answer with diagrams of valence sublevels.
4.Structural formula of nitric acid. The angles between O–N–O bonds are close to 120 o. Define:
a) type of hybridization of the nitrogen atom;
b) which AO of the nitrogen atom takes part in the formation of the -bond;
c) which AO of the nitrogen atom takes part in the formation of an -bond according to the donor-acceptor mechanism.
What do you think the angle between the H–O–N bonds in this molecule is approximately equal to? 5.Create the structural formula of the cyanide ion CN (negative charge on the carbon atom). It is known that cyanides (compounds containing such an ion) and carbon monoxide CO are strong poisons, and their biological effect is very similar. Offer your explanation of the proximity of their biological action.
7.14. Metal connection. Metals
A covalent bond is formed between atoms that are similar in their propensity to give up and gain electrons only when the sizes of the bonded atoms are small. In this case, the electron density in the region of overlapping electron clouds is significant, and the atoms turn out to be tightly bound, as, for example, in the HF molecule. If at least one of the bonded atoms has a large radius, the formation of a covalent bond becomes less advantageous, since the electron density in the region of overlapping electron clouds for large atoms is much less than for small ones. An example of such a molecule with a weaker bond is the HI molecule (using Table 21, compare the atomization energies of HF and HI molecules).
And yet between large atoms ( r o > 1.1) a chemical bond occurs, but in this case it is formed due to the sharing of all (or part) of the valence electrons of all bonded atoms. For example, in the case of sodium atoms, all 3 s-electrons of these atoms, and a single electron cloud is formed:
Atoms form a crystal with metal communication
In this way, both atoms of the same element and atoms of different elements can bond with each other. In the first case, simple substances called metals, and in the second - complex substances called intermetallic compounds.
Of all the substances with metallic bonds between atoms, you will only learn about metals in school. What is the spatial structure of metals? The metal crystal consists of atomic skeletons, remaining after the socialization of valence electrons, and the electron cloud of socialized electrons. The atomic cores usually form a very close packing, and the electron cloud occupies the entire remaining free volume of the crystal.
The main types of dense packaging are cubic closest packing(KPU) and hexagonal close packing(GPU). The names of these packages are associated with the symmetry of the crystals in which they are realized. Some metals form loosely packed crystals - body-centered cubic(OTSK). Volume and ball-and-stick models of these packages are shown in Figure 7.6.
Cubic close packing is formed by atoms of Cu, Al, Pb, Au and some other elements. Hexagonal close packing - atoms of Be, Zn, Cd, Sc and a number of others. Body-centered cubic packing of atoms is present in crystals of alkali metals, elements of VB and VIB groups. Some metals may have different structures at different temperatures. The reasons for such differences and structural features of metals are still not fully understood.
When melted, metal crystals turn into metal liquids. The type of chemical bond between atoms does not change.
The metal bond does not have directionality and saturation. In this respect it is similar to an ionic bond.
In the case of intermetallic compounds, we can also talk about the polarizability of the metallic bond.
Characteristic physical properties of metals:
1) high electrical conductivity;
2) high thermal conductivity;
3) high ductility.
The melting points of different metals are very different from each other: the lowest melting point is for mercury (- 39 o C), and the highest is for tungsten (3410 o C).
| Beryllium Be- light gray, lightweight, fairly hard, but usually brittle metal. Melting point 1287 o C. In air it becomes covered with an oxide film. Beryllium is a fairly rare metal; living organisms in the process of their evolution had practically no contact with it, so it is not surprising that it is poisonous to the animal world. It is used in nuclear technology. Zinc Zn is a white soft metal with a bluish tint. Melting point 420 o C. In air and water it is covered with a thin dense film of zinc oxide, which prevents further oxidation. In production it is used for galvanizing sheets, pipes, wires, protecting iron from corrosion. Wolfram W. It is the most refractory of all metals: the melting point of tungsten is 3387 o C. Typically, tungsten is quite brittle, but after careful cleaning it becomes ductile, which makes it possible to draw thin wire from it, from which the filaments of light bulbs are made. However, most of the tungsten produced is used for the production of hard and wear-resistant alloys that can retain these properties when heated even to 1000 o C. |
METAL, INTERMETALLIC COMPOUND, METALLIC BOND, DENSE PACKING.
1. To characterize various packages, the concept of “space filling coefficient” is used, that is, the ratio of the volume of atoms to the volume of the crystal
Where V a - volume of an atom,
Z is the number of atoms in a unit cell,
V i- volume of the unit cell.
Atoms in this case are represented by rigid balls of radius R, touching each other. Ball volume V w = (4/3) R 3 .
Determine the space filling factor for bulk and bcc packaging.
2. Using the values of metal radii (Appendix 9), calculate the unit cell size of a) copper (CPU), b) aluminum (CPU) and c) cesium (BCC).
Help me solve chemistry please. Indicate the type of bond in the molecules NH3, CaCl2, Al2O3, BaS... and received the best answer
Answer from Olga Lyabina[guru]
1) NH3 bond type cov. polar. Three unpaired electrons of nitrogen and one of hydrogen each take part in the formation of a bond. There are no pi bonds. sp3 hybridization. The shape of the molecule is pyramidal (one orbital does not participate in hybridization, the tetrahedron turns into a pyramid)
CaCl2 type of bond is ionic. The bond formation involves two calcium electrons in the s orbital, which accept two chlorine atoms, completing their third level. no pi bonds, hybridization type sp. they are located in space at an angle of 180 degrees
Al2O3 bond type is ionic. Three electrons from the s and p orbitals of aluminum are involved in the formation of the bond, which oxygen accepts, completing its second level. O=Al-O-Al=O. There are pi bonds between oxygen and aluminum. sp hybridization type most likely.
BaS type of bond is ionic. two electrons of barium are accepted by sulfur. Ba=S is one pi bond. hybridization sp. Flat molecule.
2) AgNO3
silver is reduced at the cathode
K Ag+ + e = Ag
water oxidizes at the anode
A 2H2O - 4e = O2 + 4H+
according to Faraday's law (whatever...) the mass (volume) of the substance released at the cathode is proportional to the amount of electricity passing through the solution
m(Ag) = Me/zF *I*t = 32.23 g
V(O2) = Ve/F *I*t = 1.67 l
Answer from 2 answers[guru]
Hello! Here is a selection of topics with answers to your question: Help me solve chemistry, please. Indicate the type of bond in the molecules NH3, CaCl2, Al2O3, BaS...
