How to determine pi bonds in chemistry. Single and multiple connections. Sigma and pi bonds. Covalent bonds. Pi and sigma bonds

93. Exemption from criminal liability in connection with reconciliation with the victim and in connection with a change in the situation Exemption from criminal liability in connection with reconciliation with the victim under the legislation in force before the entry into force of the 1996 Criminal Code,

7.5. Analysis of a specific situation “Holding a meeting at the Sigma company”

7.5. Analysis of a specific situation “Holding a meeting at the Sigma company” Goal: Development of skills in analyzing organizational culture using a specific example. Assignment.Review the situation below and answer the following questions.1. How do you rate the level

Pi bonds occur when p-atomic orbitals on either side of the line of atomic bonding overlap. It is believed that the pi bond is realized in multiple bonds - a double bond consists of one sigma and one pi bond, a triple bond - of one sigma and two orthogonal pi bonds.

The concept of sigma and pi bonds was developed by Linus Pauling in the 30s of the last century. One s- and three p-valence electrons of the carbon atom undergo hybridization and become four equivalent sp 3 hybridized electrons, through which four equivalent chemical bonds are formed in the methane molecule. All bonds in a methane molecule are equidistant from each other, forming a tetrahedron configuration.

In the case of double bond formation, sigma bonds are formed by sp 2 hybridized orbitals. The total number of such bonds in a carbon atom is three and they are located in the same plane. The angle between the bonds is 120°. The pi bond is located perpendicular to the indicated plane (Fig. 1).

In the case of triple bond formation, sigma bonds are formed by sp-hybridized orbitals. The total number of such bonds on a carbon atom is two and they are at an angle of 180° to each other. The two pi bonds of a triple bond are mutually perpendicular (Fig. 2).

In the case of the formation of an aromatic system, for example, benzene C 6 H 6, each of the six carbon atoms is in a state of sp 2 hybridization and forms three sigma bonds with bond angles of 120 °. The fourth p-electron of each carbon atom is oriented perpendicular to the plane of the benzene ring (Fig. 3.). In general, a single bond appears that extends to all carbon atoms of the benzene ring. Two regions of high electron density pi bonds are formed on either side of the sigma bond plane. With such a bond, all carbon atoms in the benzene molecule become equivalent and, therefore, such a system is more stable than a system with three localized double bonds. The non-localized pi bond in the benzene molecule causes an increase in the bond order between carbon atoms and a decrease in the internuclear distance, that is, the length of the chemical bond d cc in the benzene molecule is 1.39 Å, while d C-C = 1.543 Å, and d C=C = 1.353 Å.

L. Pauling's concept of sigma and pi bonds became an integral part of the theory of valence bonds. Animated images of atomic orbital hybridization have now been developed.

However, L. Pauling himself was not satisfied with the description of sigma and pi bonds. At a symposium on theoretical organic chemistry dedicated to the memory of F.A. Kekule (London, September 1958), he abandoned the σ, π-description and proposed and substantiated the theory of a bent chemical bond. The new theory clearly took into account the physical meaning of a covalent chemical bond, namely Coulomb electron correlation.

Notes

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There are two types of covalent bonds: sigma and pi bonds. A sigma bond is a single covalent bond formed when an AO overlaps along a straight line (axis) connecting the nuclei of two bonded atoms with a maximum overlap on this straight line. A sigma bond can arise when any (s-, p-hybrid) AOs overlap. In organogens (carbon, nitrogen, oxygen, sulfur), hybrid orbitals can take part in the formation of sigma bonds, providing more efficient overlap. In addition to axial overlap, another type of overlap is possible - lateral overlap of p-AO, leading to the formation of a pi bond. A pi bond is a bond formed by the lateral overlap of unhybridized p-AOs with a maximum overlap on both sides of the straight line connecting the nuclei of atoms. Multiple bonds often found in organic compounds are a combination of sigma and pi bonds; double - one sigma and one pi, triple - one sigma and two pi bonds.

Bonding energy is the energy released when a bond is formed or required to separate two bonded atoms. It serves as a measure of the strength of the bond: the greater the energy, the stronger the bond.

Bond length is the distance between the centers of bonded atoms. A double bond is shorter than a single bond, and a triple bond is shorter than a double bond. Bonds between carbon atoms in different states of hybridization are characterized by a general pattern: with an increase in the fraction of the s-orbital in the hybrid orbital, the bond length decreases. For example, in the series of compounds propane CH3-CH2-CH3, propene CH3-CH=CH2, propine CH3-C-=CH, the CH3-C bond length is respectively 0.154, 0.150 and 0.146 nm.

In chemistry, the concept of hybrid orbitals of the carbon atom and other elements is widely used. The concept of hybridization as a way to describe the rearrangement of orbitals is necessary in cases where the number of unpaired electrons in the ground state of an atom is less than the number of bonds formed. It is postulated that different atomic orbitals having similar energy levels interact with each other, forming hybrid orbitals with the same shape and energy. Hybridized orbitals, due to greater overlap, form stronger bonds compared to non-hybridized orbitals.

The type of hybridization determines the orientation of hybrid AOs in space and, consequently, the geometry of the molecules. Depending on the number of orbitals that have entered into hybridization, a carbon atom can be in one of three states of hybridization. sp3-Hybridization. As a result of sp3 hybridization, a carbon atom from the ground state 1s2-2s2-2p2 due to the movement of an electron from the 2s to 2p orbital goes into the excited state 1s2-2s1-2p3. When four external AOs of an excited carbon atom (one 2s and three 2p orbitals) are mixed, four equivalent sp-hybrid orbitals arise. They have the shape of a three-dimensional figure eight, one of the blades of which is much larger than the other. Due to mutual repulsion, sp3-hybrid AOs are directed in space towards the vertices of the tetrahedron and the angles between them are equal to 109.5° (the most favorable location). Each hybrid orbital in an atom is filled with one electron. The carbon atom in the state of sp3 hybridization has the electronic configuration 1s2(2sp3)4.

This state of hybridization is characteristic of carbon atoms in saturated hydrocarbons (alkanes) and, accordingly, in alkyl radicals of their derivatives. sp2-Hybridization. As a result of sp2 hybridization, due to the mixing of one 2s and two 2p AOs of an excited carbon atom, three equivalent sp2 hybrid orbitals are formed, located in the same plane at an angle of 120’. Unhybridized 2p-AO is in a perpendicular plane. The carbon atom in the state of sp2 hybridization has the electronic configuration 1s2-(2sp2)3-2p1. This carbon atom is characteristic of unsaturated hydrocarbons (alkenes), as well as some functional groups, for example carbonyl, carboxyl, etc. sp-Hybridization. As a result of sp hybridization, due to the mixing of one 2s and one 2p orbitals of an excited carbon atom, two equivalent sp hybrid AOs are formed, located linearly at an angle of 180°. The two remaining unhybridized 2p-AOs are located in mutually perpendicular planes. The carbon atom in the state of sp-hybridization has the electronic configuration 1s2-(2sp)2-2p2. Such an atom is found in compounds that have a triple bond, for example in alkynes and nitriles. Atoms of other elements can also be in a hybridized state. For example, the nitrogen atom in the ammonium ion NH4+ and, accordingly, alkylammonium RNH3+ is in a state of sp3 hybridization; in pyrrole and pyridine - sp2-hybridization; in nitriles - sp-hybridization.

Single connection– a covalent bond in which only one shared electron pair is formed between two atoms.

Sigma communication– a covalent bond, during the formation of which the area of ​​overlap of electron clouds is located on the line connecting the nuclei of atoms. Single bonds are always sigma bonds.

Pi connection– a covalent bond, during the formation of which the area of ​​overlapping electron clouds is located on both sides of the line connecting the nuclei. They are formed when two or three common electron pairs appear between two atoms. The number of shared electron pairs between bonded atoms characterizes communication multiplicity.

If a bond between two atoms is formed by two shared electron pairs, then such a bond is called double bond. Any double bond consists of one sigma bond and one pi bond.

If a bond between two atoms is formed by three shared electron pairs, then such a bond is called triple bond. Any triple bond consists of one sigma bond and two pi bonds.

Double and triple bonds have a common name: multiple bonds.

The overlapping orbitals must have the same symmetry about the internuclear axis. The overlap of atomic orbitals along the line connecting the atomic nuclei leads to the formation of σ - connections. Only one σ bond is possible between two atoms in a chemical particle. All σ bonds have axial symmetry relative to the internuclear axis. Fragments of chemical particles can rotate around the internuclear axis without disturbing the degree of overlap of atomic orbitals forming σ bonds. A set of directed, strictly oriented in space σ-bonds creates the structure of a chemical particle.
With the additional overlap of atomic orbitals perpendicular to the bond line, π bonds. As a result, multiple bonds arise between atoms: Single (σ), Double (σ +π), Triple (σ + π + π).F−F, O=O, N≡N.
With the appearance of a π-bond that does not have axial symmetry, free rotation of fragments of a chemical particle around the σ-bond becomes impossible, since it should lead to the rupture of the π-bond. In addition to σ- and π-bonds, it is possible to form another type of bond - δ-bonds: Typically, such a bond is formed after the formation of σ- and π-bonds by atoms when the atoms have d- and f-orbitals by overlapping their “petals” in four places at once. As a result, the multiplicity of communication can increase to 4-5.

Main types of structures of inorganic compounds. Substances with molecular and
non-molecular structure. Atomic, molecular, ionic and metallic
crystal lattices.

Molecular and non-molecular substances - one of the characteristics of chemical substances regarding their structure.

Molecular substances- these are substances whose smallest structural particles are molecules

Molecules- the smallest particle of a molecular substance that can exist independently and retain its chemical properties. Molecular substances have low melting and boiling points and exist under standard conditions in a solid, liquid or gaseous state.

Non-molecular substances- these are substances whose smallest structural particles are atoms or ions.

And he is an atom or group of atoms that has a positive or negative charge.

Non-molecular substances are in standard conditions in a solid state of aggregation and have high melting and boiling points.

There are substances with molecular and non-molecular structure. All gases and all liquids have a molecular structure. Solids can have a molecular or non-molecular structure. Volatile solids (ice, iodine, white phosphorus, organic substances) have a molecular structure. Molecules are located at the nodes of the crystal lattice of highly volatile solids. Most inorganic solids have a non-molecular structure; the lattice sites contain ions (salts, bases) or atoms (metals, diamond, silicon). Substances with a molecular structure make up more than 95% of all known substances, since organic substances have a molecular structure, and much more organic substances are known than inorganic ones.
Chemical reactions. Classification of chemical reactions. The main problems of chemical kinetics and chemical thermodynamics.

Chemical reactions– These are phenomena in which the transformation of one substance into another occurs.

Signs of chemical reactions:

ü Gas release

Na 2 CO 3 +2HCl=2NaCl+H 2 O+CO 2

ü Precipitation or dissolution of sediment

BaCl 2 +H 2 SO 4 =BaSO 4 +2HCl

ü Color change

FeCl 3(yellow) +3KSCN (colorless) =Fe(SCN) 3(red) +3KCl

ü Odor appears.

ü Emission of light and heat

H 2 SO 4 +2NaOH=Na 2 SO 4 +2H 2 O+Q

2Mg+O 2 =2MgO+ hv

For chemical reactions to occur, conditions are necessary: ​​contact of reacting substances, heating, lighting.

Classifications of chemical reactions:

I. According to the number and composition of the starting reagents:

a) Compound reaction- a reaction in which several substances form one substance, more complex than the original ones: A+B=AB

SO 3 +H 2 O=H 2 SO 4

NH 3 +HCl=NH 4 Cl

b) Decomposition reaction- a reaction in which several substances are formed from one complex substance. The final products can be both simple and complex substances: AB=A+B

2KClO 3 =2KCl+3O 2

c) Substitution reaction- a reaction in which atoms of one element replace atoms of another element in a complex substance and at the same time two new ones are formed - simple and complex: X+AB=AX+B

Fe+CuSO 4 =FeSO 4 +Cu

Zn+2HCl=ZnCl 2 +H 2

d) Exchange reaction- a reaction in which reacting substances exchange their constituent parts, as a result of which two new complex substances are formed from two complex substances: AB+CX=AX+CB

BaCl 2 + Na 2 SO 4 = 2NaCl + BaSO 4

AgNO 3 +HCl=HNO 3 +AgCl

II. According to the sign of the thermal effect, reactions are divided into:

a) endothermic- reactions that occur with the absorption of heat

b) exothermic- reactions that result in the release of heat.

III. Taking into account the phenomenon of catalysis:

a) catalytic(flowing with the participation of a catalyst)

b) non-catalytic.

IV. Based on reversibility, reactions are divided into:

a) reversible– flowing simultaneously in the forward and reverse directions

b) irreversible – flowing in one direction

V. Based on changes in the oxidation states of elements in the molecules of reacting substances:

a) OVR– electron transfer reactions

b) Not OVR– reactions without electron transfer.

VI. Based on the homogeneity of the reaction system:

a) Homogeneous– flowing in a homogeneous system

b) Heterogeneous– occurring in a heterogeneous system

14. Basic characteristics of covalent bonds. Bond length and energy. Saturation and direction. Multiplicity of communication. Sigma and pi connections.

- A chemical bond carried out by shared electron pairs is called atomic or covalent. Each covalent chemical bond has certain qualitative or quantitative characteristics. These include:

Link length

Communication energy

Saturability

Communication direction

Communication polarity

Communication multiplicity

- Link length– the distance between the nuclei of bonded atoms. It depends on the size of the atoms and the degree of overlap of their electron shells. The length of a bond is determined by the order of the bond: the higher the order of the bond, the shorter its length.

Communication energy is the energy that is released when a molecule is formed from single atoms. It is usually expressed in J/mol (or cal/mol). Bond energy is determined by the bond order: the greater the bond order, the greater its energy. Bond energy is a measure of its strength. Its value is determined by the work required to break a bond, or the gain in energy when a substance is formed from individual atoms. The system that contains less energy is more stable. For diatomic molecules, the bond energy is equal to the dissociation energy taken with the opposite sign. If more than 2 different atoms combine in a molecule, then the average binding energy does not coincide with the dissociation energy of the molecule. Bond energies in molecules consisting of identical atoms decrease in groups from top to bottom. The bond energies increase over the period.

- Saturability– shows how many bonds a given atom can form with others due to shared electron pairs. It is equal to the number of common electron pairs with which a given atom is connected to others. The saturation of a covalent bond is the ability of an atom to participate in the formation of a limited number of covalent bonds.

Focus– this is a certain relative arrangement of connecting electron clouds. It leads to a certain arrangement in space of the nuclei of chemically bonded atoms. The spatial orientation of a covalent bond is characterized by the angles between the bonds formed, which are called bond angles.

- Multiplicity of communication. Determined by the number of electron pairs involved in the bond between atoms. If a bond is formed by more than one pair of electrons, then it is called multiple. As the bond multiplicity increases, the energy increases and the bond length decreases. In molecules with a multiple bond there is no rotation around an axis.

- Sigma and pi bonds. The chemical bond is caused by the overlap of electron clouds. If this overlap occurs along a line connecting the atomic nuclei, then the bond is called a sigma bond. It can be formed by s-s electrons, p-p electrons, s-p electrons. A chemical bond carried out by one electron pair is called a single bond. Single bonds are always sigma bonds. S-type orbitals form only sigma bonds. But a large number of compounds are known that have double and even triple bonds. One of them is sigma bond and the others are called pi bonds. When such bonds are formed, overlapping electron clouds occur in two regions of space symmetrical to the internuclear axis.

15. Hybridization of atomic orbitals using the example of molecules: methane, aluminum chloride, beryllium chloride. Bond angle and molecular geometry. Molecular orbital method (MO LCAO). Energy diagrams of homo- and hetero-nuclear molecules (N2, Cl2, N.H.3, Be2).

- Hybridization. The new set of mixed orbitals is called hybrid orbitals, and the mixing technique itself is called hybridization of atomic orbitals.

The mixing of one s and one p orbital, as in BeCl2, is called sp hybridization. In principle, hybridization of an s-orbital is possible not only with one, but also with two, three, or a non-integer number of p-orbitals, as well as hybridization involving d-orbitals.

Let's consider the linear BeCl2 molecule. A beryllium atom in the valence state is capable of forming two bonds due to one s- and one p-electron. Obviously, this should result in two bonds with chlorine atoms of different lengths, since the radial distribution of these electrons is different. The real BeCl2 molecule is symmetrical and linear; its two Be-Cl bonds are exactly the same. This means that they are provided with electrons of the same state, i.e. here the beryllium atom in the valence state no longer has one s- and one p-electron, but two electrons located in orbitals formed by the “mixing” of s- and p-atomic orbitals. A methane molecule will have sp3 hybridization, and an aluminum chloride molecule will have sp2 hybridization.

Conditions for hybridization stability:

1) Compared to the original orbital atoms, the hybrid orbitals should overlap more closely.

2) Atomic orbitals that are close in energy level take part in hybridization; therefore, stable hybrid orbitals should be formed on the left side of the periodic table.

- MO LCAO. Molecular orbitals can be thought of as a linear combination of atomic orbitals. Molecular orbitals must have a certain symmetry. When filling atomic orbitals with electrons, it is necessary to take into account the following rules:

1. If an atomic orbital is a certain function that is a solution to the Schrödinger Equation and describes the state of an electron in an atom, the MO method is also a solution to the Schrödinger equation, but for an electron in a molecule.

2. A molecular orbital is found by adding or subtracting atomic orbitals.

3. Molecular orbitals and their number are equal to the sum of the atomic orbitals of the reacting atoms.

If the solution for molecular orbitals is obtained by adding the functions of atomic orbitals, then the energy of the molecular orbitals will be lower than the energy of the original atomic orbitals. And such an orbital is called bonding orbital.

In the case of subtraction of functions, the molecular orbital has a higher energy, and it is called loosening.

There are sigma and pi orbitals. They are filled in according to Hund's rule.

The number of bonds (bond order) is equal to the difference between the total number of electrons in the bonding orbital and the number of electrons in the antibonding orbital, divided by 2.

The MO method uses energy diagrams:

16. Polarization of communication. Dipole moment of connection. Characteristics of interacting atoms: ionization potential, electron affinity, electronegativity. The degree of ionicity of the bond.

- Dipole moment- a physical quantity characterizing the electrical properties of a system of charged particles. In the case of a dipole (two particles with opposite charges), the electric dipole moment is equal to the product of the positive charge of the dipole and the distance between the charges and is directed from the negative charge to the positive one. The dipole moment of a chemical bond is caused by the displacement of the electron cloud towards one of the atoms. A bond is called polar if the corresponding dipole moment differs significantly from zero. There are cases when individual bonds in a molecule are polar, and the total dipole moment of the molecule is zero; such molecules are called non-polar (for example, CO 2 and CCl 4 molecules). If the dipole moment of a molecule is non-zero, the molecule is called polar. For example, the H 2 O molecule. The order of magnitude of the dipole moment of the molecule is determined by the product of the electron charge (1.6.10 -19 C) and the length of the chemical bond (about 10 -10 m).

The chemical nature of an element is determined by the ability of its atom to lose and gain electrons. This ability can be quantified by the ionization energy of an atom and its electron affinity.

- Ionization energy of an atom is the amount of energy required to remove an electron from an unexcited atom. It is expressed in kilojoules per mole. For multi-electron atoms, the ionization energies E1, E2, E3, ..., En correspond to the separation of the first, second, etc. electrons. In this case, always E1

- Atom electron affinity– the energetic effect of adding an electron to a neutral atom, transforming it into a negative ion. The electron affinity of an atom is expressed in kJ/mol. Electron affinity is numerically equal but opposite in sign to the ionization energy of a negatively charged ion and depends on the electronic configuration of the atom. Group 7 p-elements have the highest electron affinity. Atoms with the s2 (Be, Mg, Ca) and s2p6 (Ne, Ar, Kr) configuration or half-filled with a p-sublayer (N, P, As) do not exhibit electron affinity.

- Electronegativity- an averaged characteristic of the ability of an atom in a compound to attract an electron. In this case, the difference in the states of atoms in different compounds is neglected. Unlike ionization potential and electron affinity, EO is not a strictly defined physical quantity, but a useful conditional characteristic. The most electronegative element is fluorine. EO depends on ionization energy and electron affinity. According to one definition, the EO of an atom can be expressed as half the sum of its ionization energy and electron affinity. An element cannot be assigned a constant EO. It depends on many factors, in particular on the valence state of the element, the type of compound in which it is included, etc.

17. Polarizing ability and polarizing effect. Explanation of some physical properties of substances from the point of view of this theory.

- Polarization theory considers all substances to be purely ionic. In the absence of an external field, all ions have a spherical shape. When the ions approach each other, the field of the cation affects the field of the anion, and they are deformed. Ion polarization is the displacement of the outer electron cloud of ions relative to their nucleus.

Polarization consists of two processes:

ion polarizability

polarizing effect on another ion

The polarizability of an ion is a measure of the ability of the ion's electron cloud to deform under the influence of an external electric field.

Regularities of ion polarizability:

Anions are more polarized than cations. Excessive electron density leads to high diffuseness and looseness of the electron cloud.

The polarizability of isoelectronic ions increases with a decrease in positive and increasing negative charges. Isoelectronic ions have the same configuration.

In multiply charged cations, the nuclear charge is much greater than the number of electrons. This compacts the electron shell and stabilizes it, so such ions are less susceptible to deformation. The polarizability of cations decreases upon transition from ions with an outer electron shell filled with 18 electrons to an unfilled one, and then to noble gas ions. This is due to the fact that for electrons of the same period, the d-electron shell is more diffuse compared to the s- and p-electron shells, because d electrons spend more time near the nucleus. Therefore, d-electrons interact more strongly with surrounding anions.

The polarizability of analogue ions increases with increasing number of electronic layers. Polarizability is most difficult for small-sized and multiply charged cations, with an electron shell of noble gases. Such cations are called hard. Bulk multicharged anions and low-charge bulk cations are most easily polarized. These are soft ions.

- Polarizing effect. Depends on the charges, size and structure of the outer electronic layer.

1. The polarizing effect of a cation increases with increasing its charge and decreasing radius. The maximum polarizing effect is characteristic of Katons with small radii and large charges, therefore they form covalent-type compounds. The greater the charge, the greater the polarizing bond.

2. The polarizing effect of cations increases with the transition from ions with an s-electron cloud to an incomplete one and to an 18-electron one. The greater the polarizing effect of the cation, the greater the contribution of the covalent bond.

- Application of polarization theory to explain physical properties:

The greater the polarizability of an anion (the polarizing effect of a cation), the more likely it is to form a covalent bond. Therefore, the boiling and melting points of compounds with covalent bonds will be lower than those with ionic bonds. The higher the ionicity of the bond, the higher the melting and boiling points.

The deformation of the electron shell affects the ability to reflect or absorb light waves. From here, from the perspective of the theory of polarization, the color of compounds can be explained: white reflects everything; black – absorbs; transparent – ​​lets through. This is related: if the shell is deformed, then the quantum levels of the electrons move closer together, reducing the energy barrier, so low energy is required for excitation. Because absorption is associated with the excitation of electrons, i.e. with their transition to high-lying levels, then in the presence of high polarization, already visible light can excite external electrons and the substance will be colored. The higher the charge of the anion, the lower the color intensity. The polarizing effect affects the reactivity of compounds; therefore, for many compounds, salts of oxygen-containing acids are more stable than the salts themselves. The greatest polarizing effect is found in d-elements. The greater the charge, the greater the polarizing effect.

18. Ionic bond as a limiting case of a covalent polar bond. Properties of substances with different types of bonds.

The nature of the ionic bond can be explained by the electrostatic interaction of ions. The ability of elements to form simple ions is determined by the structure of their atoms. Cations most easily form elements with low ionization energy, alkali and alkaline earth metals. Anions are most easily formed by p-elements of group 7, due to their high electron affinity.

The electrical charges of ions cause their attraction and repulsion. Ions can be thought of as charged balls whose force fields are uniformly distributed in all directions in space. Therefore, each ion can attract ions of the opposite sign to itself in any direction. An ionic bond, unlike a covalent bond, is characterized by non-directionality.

The interaction of ions of opposite signs with each other cannot lead to complete mutual compensation of their force fields. Because of this, they retain the ability to attract ions in other directions. Therefore, unlike a covalent bond, an ionic bond is characterized by unsaturation.

19.Metal connection. Similarities and differences with ionic and covalent bonds

A metallic bond is a bond in which the electrons of each individual atom belong to all the atoms in contact. The energy difference between “molecular” orbitals in such a bond is small, so electrons can easily move from one “molecular” orbital to another and, therefore, move in the volume of the metal.

Metals differ from other substances in their high electrical and thermal conductivity. Under normal conditions, they are crystalline substances (with the exception of mercury) with high coordination numbers of atoms. In a metal, the number of electrons is much less than the number of orbitals, so electrons can move from one orbital to another. Metal atoms are characterized by high ionization energy - valence electrons are weakly retained in the atom, i.e. easily move in the crystal. The ability of electrons to move around a crystal determines the electrical conductivity of metals.

Thus, unlike covalent and ionic compounds, in metals a large number of electrons simultaneously bind a large number of atomic nuclei, and the electrons themselves can move in the metal. In other words, in metals there is a highly delocalized chemical bond. The metallic bond has a certain similarity to the covalent bond, since it is based on the sharing of valence electrons. However, in the formation of a covalent bond, the valence electrons of only two interacting atoms participate, while in the formation of a metallic bond, all atoms take part in the sharing of electrons. That is why the metallic bond does not have spatial directionality and saturation, which largely determines the specific properties of metals. The energy of a metallic bond is 3-4 times less than the energy of a covalent bond.

20. Hydrogen bond. Intermolecular and intramolecular. Mechanism of education. Features of the physical properties of substances with hydrogen bonds. Examples.

- A hydrogen bond is a special type of chemical bond. It is characteristic of hydrogen compounds with the most electronegative elements (fluorine, oxygen, nitrogen and, to a lesser extent, chlorine and sulfur).

Hydrogen bonding is very common and plays an important role in the association of molecules, in the processes of crystallization, dissolution, formation of crystal hydrates, etc. For example, in the solid, liquid and even gaseous state, hydrogen fluoride molecules are connected in a zigzag chain, which is due precisely to the hydrogen bond.

Its peculiarity is that a hydrogen atom, which is part of one molecule, forms a second, weaker bond with an atom in another molecule, as a result of which both molecules are combined into a complex. A characteristic feature of such a complex is the so-called hydrogen bridge – A – H...B–. The distance between atoms in a bridge is greater than between atoms in a molecule. Initially, hydrogen bonding was interpreted as an electrostatic interaction. It has now been concluded that the donor-acceptor interaction plays a major role in hydrogen bonding. Hydrogen bonds are formed not only between molecules of different substances, but also in molecules of the same substance H2O, HF, NH3, etc. This also explains the difference in the properties of these substances compared to related compounds. Hydrogen bonding within molecules is known, especially in organic compounds. Its formation is facilitated by the presence in the molecule of the acceptor group A-H and the donor group B-R. In the A-H molecule, A is the most electronegative element. The formation of hydrogen bonding in polymers, such as peptides, results in a helical structure. DNA, deoxyribonucleic acid, the keeper of the code of heredity, has similar structures. Hydrogen bonds are not strong. They easily form and break at ordinary temperatures, which is very important in biological processes. It is known that hydrogen compounds with highly electronegative nonmetals have abnormally high boiling points.

Intermolecular interaction. The forces of attraction between saturated atoms and molecules are extremely weak compared to ionic and covalent bonds. Substances in which molecules are held together by extremely weak forces are often gases at 20 degrees, and in many cases their boiling points are very low. The existence of such weak forces was discovered by Van der Waals. The existence of such forces in the system can be explained:

1. The presence of a permanent dipole in the molecule. In this case, as a result of the simple electrostatic attraction of dipoles, weak interaction forces arise - dipole-dipole (H2O, HCl, CO)

2. The dipole moment is very small, but when interacting with water, an induced dipole can be formed, which arises as a result of the polymerization of molecules by the dipoles of surrounding molecules. This effect can be superimposed on the dipole-dipole interaction and increase the attraction.

3. Dispersion forces. These forces act between any atoms and molecules, regardless of their structure. London introduced this concept. For symmetric atoms, the only forces acting are the London forces.

21. Aggregate states of matter: solid, liquid, gaseous. Crystalline and amorphous states. Crystal lattices.

- Under ordinary conditions, atoms, ions and molecules do not exist individually. It always consists only of parts of a higher organization of a substance that practically participates in chemical transformations - the so-called state of aggregation. Depending on external conditions, all substances can be in different states of aggregation - gas, liquid, solid. The transition from one state of aggregation to another is not accompanied by a change in the stoichiometric composition of the substance, but is necessarily associated with a greater or lesser change in its structure.

Solid state- this is a state in which a substance has its own volume and its own shape. In solids, the interaction forces between particles are very strong. Almost all substances exist in the form of several solids. The reactivity and other properties of these bodies are usually different. The ideal solid state corresponds to a hypothetical ideal crystal.

Liquid state- this is a state in which a substance has its own volume, but does not have its own shape. The liquid has a certain structure. In structure, the liquid state is intermediate between a solid state with a strictly defined periodic structure and a gas in which there is no structure. Hence, a liquid is characterized, on the one hand, by the presence of a certain volume, and on the other, by the absence of a certain shape. The continuous movement of particles in a liquid determines strongly pronounced self-diffusion and its fluidity. The structure and physical properties of a liquid depend on the chemical identity of the particles that form it.

Gaseous state. A characteristic feature of the gas state is that the molecules (atoms) of the gas are not held together, but move freely in the volume. Intermolecular interaction forces occur when molecules come close to each other. Weak intermolecular interaction determines the low density of gases and their main characteristic properties - the desire for infinite expansion and the ability to exert pressure on the walls of vessels that impede this desire. Due to the weak intermolecular interaction at low pressure and high temperatures, all typical gases behave approximately the same, but already at ordinary temperatures and pressure the individuality of gases begins to appear. The state of a gas is characterized by its temperature, pressure and volume. The gas is considered to be at no. if its temperature is 0 degrees and the pressure is 1* 10 Pa.

- Crystalline state. Among solids, the main one is the crystalline state, characterized by a certain orientation of particles (atoms, ions, molecules) relative to each other. This also determines the external form of the substance in the form of crystals. Single crystals - single crystals exist in nature, but they can be obtained artificially. But most often crystalline bodies are polycrystalline formations - these are intergrowths of a large number of small crystals. A characteristic feature of crystalline bodies, resulting from their structure, is anisotropy. It manifests itself in the fact that the mechanical, electrical and other properties of crystals depend on the direction of the external influence of forces on the crystal. Particles in crystals undergo thermal vibrations around the equilibrium position or around the nodes of the crystal lattice.

Amorphous state. The amorphous state is similar to the liquid state. It is characterized by incomplete ordering of the relative arrangement of particles. The bonds between structural units are not equivalent, therefore amorphous bodies do not have a specific melting point - during the heating process they gradually soften and melt. For example, the temperature range of melting processes for silicate glasses is 200 degrees. In amorphous bodies, the nature of the arrangement of atoms practically does not change when heated. Only the mobility of atoms changes - their vibrations increase.

- Crystal lattices:

Crystal lattices can be ionic, atomic (covalent or metallic) and molecular.

The ionic lattice consists of ions of opposite signs alternating at the sites.

In atomic lattices, atoms are connected by covalent or metallic bonds. Example: diamond (atomic-covalent lattice), metals and their alloys (atomic-metal lattice). The nodes of a molecular crystal lattice are formed by molecules. In crystals, molecules are connected through intermolecular interactions.

Differences in the type of chemical bond in crystals determine significant differences in the type of physical and chemical properties of a substance with all types of crystal lattice. For example, substances with an atomic-covalent lattice are characterized by high hardness, and those with an atomic-metal lattice are characterized by high plasticity. Substances with an ionic lattice have a high melting point and are not volatile. Substances with a molecular lattice (intermolecular forces are weak) are fusible, volatile, and their hardness is not high.

22. Complex compounds. Definition. Compound.

Complex compounds are molecular compounds, the combination of components of which leads to the formation of complex ions capable of free existence, both in a crystal and in solution. Complex ions are the result of interactions between the central atom (complexing agent) and the surrounding ligands. Ligands are both ions and neutral molecules. Most often, the complexing agent is a metal, which, together with ligands, forms the inner sphere. There is an outer sphere. The inner and outer spheres are interconnected by an ionic bond.