The most important characteristics of a bond include: length, polarity, dipole moment, saturation, directionality, strength, and multiplicity of the bond.

Link length– is the distance between the nuclei of atoms in a molecule. The bond length is determined by the size of the nuclei and the degree of overlap of the electron clouds.

The bond length in HF is 0.92∙10 -10, in HCl – 1.28∙10 -10 m. The shorter its length, the stronger the chemical bond.

Bond angle (Bond angle) call the angle between imaginary lines passing through the nuclei of chemically bonded atoms. ∟HOH=104 0 .5; ∟H 2 S=92.2 0; ∟H 2 S e =91 0 .0.

The most important characteristic chemical bond is energy, defining it strength.

The strength of a bond is quantitatively characterized by the energy expended to break it, and is measured in kJ per 1 mole of substance.

Therefore, the bond strength is quantitatively characterized by the sublimation energy E subl. substances and energy of dissociation of a molecule into atoms E diss. . Sublimation energy refers to the energy expended to transition a substance from a solid to a gaseous state. For diatomic molecules, the binding energy is equal to the energy of dissociation of the molecule into two atoms.

For example, E diss. (and therefore E St.) in the H 2 molecule is 435 kJ/mol. In the F 2 molecule = 159 kJ/mol, in the N 2 molecule = 940 kJ/mol.

For not diatomic, but polyatomic molecules of type AB n, the average binding energy

by AB n =A+nB.

For example, the energy absorbed during the process

equal to 924 kJ/mol.

Communication energy

E OH = = = = 462 kJ/mol.

A conclusion about the structure of molecules and the structure of a substance is made based on the results obtained different methods. In this case, the obtained information is used not only about bond lengths and energies, bond angles, but also other properties of the substance, such as magnetic, optical, electrical, thermal and others.

The set of experimentally obtained data on the structure of matter complements and generalizes the results of quantum chemical calculation methods that use the concept of the quantum mechanical theory of chemical bonding. Chemical bonding is believed to be primarily mediated by valence electrons. For s- and p-elements, the valence electrons are the electrons of the orbitals of the outer layer, and for d-elements, the electrons are the s-orbitals of the outer layer and the d-orbitals of the pre-outer layer.

The nature of the chemical bond.

A chemical bond is formed only if, as atoms approach each other, the total energy of the system (E kin. + E pot.) decreases.

Let's consider the nature of a chemical bond using the example of the molecular hydrogen ion H 2 +. (It is obtained by irradiating hydrogen molecules with H 2 electrons; in a gas discharge). For such a simple molecular system, the Schrödinger equation is most accurately solved.

In the hydrogen ion H 2 + one electron moves in the field of two nuclei - protons. The distance between the nuclei is 0.106 nm, the binding energy (dissociation into H atoms and H + ion) is 255.7 kJ/mol. That is, the particle is durable.

In the molecular ion H 2 + there are two types of electrostatic forces - the force of attraction of an electron to both nuclei and the force of repulsion between nuclei. The repulsive force manifests itself between the positively charged nuclei H A + and H A +, which can be represented in the form of the following figure. 3. The repulsive force tends to push the nuclei apart from each other.

Rice. 3. The force of repulsion (a) and attraction (b) between two nuclei, which arises when they approach each other at distances on the order of the size of atoms.

Attractive forces act between the negatively charged electron e - and the positively charged nuclei H + and H +. A molecule is formed if the resultant of the forces of attraction and repulsion is zero, that is, the mutual repulsion of nuclei must be compensated by the attraction of the electron to the nuclei. Such compensation depends on the location of the electron e - relative to the nuclei (Fig. 3 b and c). What is meant here is not the position of the electron in space (which cannot be determined), but the probability of finding the electron in space. Location of electron density in space, corresponding to Fig. 3.b) promotes the convergence of nuclei, and the corresponding Fig. 3.c) – repulsion of nuclei, since in this case the attractive forces are directed in one direction and the repulsion of nuclei is not compensated. Thus, there is a binding region, when the electron density is distributed between the nuclei, and an antibonding or antibonding region, when the electron density is distributed behind the nuclei.

If an electron enters the bonding region, a chemical bond is formed. If the electron falls into the antibonding region, then a chemical bond is not formed.

Depending on the nature of the electron density distribution in the binding region, three main types of chemical bonds are distinguished: covalent, ionic and metallic. IN pure form these bonds do not take place, and there is usually a combination of these bond types present in connections.

Types of connections.

In chemistry, the following types of bonds are distinguished: covalent, ionic, metallic, hydrogen bond, van der Waals bond, donor-acceptor bond, dative bond.

Covalent bond

When a covalent bond is formed, atoms share electrons with each other. An example of a covalent bond is the chemical bond in the Cl 2 molecule. Lewis (1916) first proposed that in such a bond, each of the two chlorine atoms shares one of its outer electrons with the other chlorine atom. To overlap atomic orbitals, two atoms must come as close to each other as possible. A shared pair of electrons forms a covalent bond. These electrons occupy the same orbital, and their spins are directed in opposite directions.

Thus, a covalent bond is accomplished by sharing electrons from different atoms as a result of pairing of electrons with opposite spins.

Covalent bonding is a common type of bonding. Covalent bonds can occur not only in molecules, but also in crystals. It occurs between identical atoms (in molecules of H 2, Cl 2, diamond) and between different atoms (in molecules of H 2 O, NH 3 ...)

Mechanism of covalent bond formation

Let us consider the mechanism using the example of the formation of the H 2 molecule.

H+H=H 2, ∆H=-436 kJ/mol

The nucleus of a free hydrogen atom is surrounded by a spherically symmetric electron cloud formed by a 1s electron. When atoms approach a certain distance, their electron clouds (orbitals) partially overlap (Fig. 4).

Rice. 4. The mechanism of bond formation in a hydrogen molecule.

If the hydrogen atoms approaching before touching have a distance between the nuclei of 0.106 nm, then after the electron clouds overlap, this distance is 0.074 nm.

As a result, a molecular two-electron cloud appears between the centers of the nuclei, which has a maximum electron density in the space between the nuclei. An increase in the negative charge density between nuclei favors a strong increase in the attractive forces between nuclei, which leads to the release of energy. The greater the overlap of electron orbitals, the stronger the chemical bond. As a result of the formation of a chemical bond between two hydrogen atoms, each of them reaches the electronic configuration of a noble gas atom - helium.

There are two methods that explain from a quantum mechanical point of view the formation of the area of ​​overlap of electron clouds, and the formation of a covalent bond, respectively. One of them is called the BC (valence bonds) method, the other MO (molecular orbitals).

The valence bond method considers the overlap of atomic orbitals of a selected pair of atoms. In the MO method, the molecule is considered as a whole and the distribution of electron density (from one electron) is spread over the entire molecule. From the position of MO 2H in H 2 are connected due to the attraction of nuclei to the electron cloud located between these nuclei.

Illustration of a covalent bond

Connections are depicted in different ways:

1). Using electrons as dots

In this case, the formation of a hydrogen molecule is shown by the diagram

N∙ + N∙ → N: N

2). Using square cells (orbitals), like placing two electrons with opposite spins in one molecular quantum cell

This diagram shows that the molecular energy level is lower than the original atomic levels, which means the molecular state of the substance is more stable than the atomic one.

3). A covalent bond is represented by a line

For example, H – N. This line symbolizes a pair of electrons.

If one covalent bond (one common electron pair) occurs between atoms, then it is called single, if more, then a multiple double(two common electron pairs), triple(three common electron pairs). A single bond is represented by one line, a double bond by two lines, and a triple bond by three lines.

The dash between the atoms shows that they have a generalized pair of electrons.

Classification of covalent bonds

Depending on the direction of overlap of electron clouds, σ-, π-, δ-bonds are distinguished. The σ bond occurs when electron clouds overlap along the axis connecting the nuclei of interacting atoms.

Examples of σ-bonds:

Rice. 5. Formation of a σ bond between s-, p-, d- electrons.

An example of the formation of a σ bond when s-s clouds overlap is observed in the hydrogen molecule.

The π bond occurs when the electron clouds on either side of the axis overlap, connecting the nuclei of atoms.

Rice. 6. Formation of π-bond between p-, d- electrons.

δ-coupling occurs when two d-electron clouds located in parallel planes overlap. The δ bond is less strong than the π bond, and the π bond is less strong than the σ bond.

Properties of covalent bonds

A). Polarity.

There are two types of covalent bonds: nonpolar and polar.

In the case of a nonpolar covalent bond, the electron cloud formed by a common pair of electrons is distributed in space symmetrically relative to the atomic nuclei. An example is diatomic molecules consisting of atoms of one element: H 2, Cl 2, O 2, N 2, F 2. Their electron pair belongs equally to both atoms.

In the case of a polar bond, the electron cloud forming the bond is shifted toward the atom with higher relative electronegativity.

Examples are the following molecules: HCl, H 2 O, H 2 S, N 2 S, NH 3, etc. Consider the formation of an HCl molecule, which can be represented by the following diagram

The electron pair is shifted to the chlorine atom, because the relative electronegativity of the chlorine atom (2.83) is greater than that of the hydrogen atom (2.1).

b). Saturability.

The ability of atoms to participate in the formation of a limited number of covalent bonds is called the saturation of a covalent bond. The saturation of covalent bonds is due to the fact that only electrons from external energy levels, that is, a limited number of electrons, participate in chemical interactions.

V) . Focus and covalent bond hybridization.

A covalent bond is characterized by directionality in space. This is explained by the fact that electron clouds have a certain shape and their maximum overlap is possible at a certain spatial orientation.

The direction of a covalent bond determines the geometric structure of molecules.

For example, for water it has a triangular shape.

Rice. 7. Spatial structure of a water molecule.

It has been experimentally established that in a water molecule H 2 O the distance between the hydrogen and oxygen nuclei is 0.096 nm (96 pm). The angle between the lines passing through the nuclei is 104.5 0. Thus, the water molecule has an angular shape and its structure can be expressed in the form of the presented figure.

Hybridization

As experimental and theoretical research(Slater, Pauling) during the formation of some compounds, such as BeCl 2, BeF 2, BeBr 2, the state of the valence electrons of an atom in a molecule is described not by pure s-, p-, d- wave functions, but by their linear combinations. Such mixed structures are called hybrid orbitals, and the mixing process is called hybridization.

As quantum chemical calculations show, mixing the s- and p-orbitals of an atom is a process favorable for the formation of a molecule. In this case, more energy is released than in the formation of bonds involving pure s- and p-orbitals. Therefore, the hybridization of the electronic orbitals of an atom leads to a large decrease in the energy of the system and, accordingly, an increase in the stability of the molecule. The hybridized orbital is more elongated on one side of the nucleus than on the other. Therefore, the electron density in the region of overlap of the hybrid cloud will be greater than the electron density in the region of overlap of the s- and p-orbitals separately, as a result of which the bond formed by the electrons of the hybrid orbital is characterized by greater strength.

Several types of hybrid states occur. When s- and p-orbitals hybridize (called sp-hybridization), two hybrid orbitals arise, located at an angle of 180 0 relative to each other. In this case, it is formed linear structure. This configuration (structure) is known for most alkaline earth metal halides (for example, BeX 2, where X = Cl, F, Br), i.e. The bond angle is 180 0 C.

Rice. 8. sp hybridization

Another type of hybridization, called sp 2 hybridization (formed from one s and two p orbitals), leads to the formation of three hybrid orbitals, which are located at an angle of 120 0 to each other. In this case, a trigonal structure of the molecule (or a regular triangle) is formed in space. Such structures are known for compounds BX 3 (X=Cl, F, Br).

Rice. 9. sp 2 -hybridization.

No less common is sp 3 hybridization, which is formed from one s- and three p- orbitals. In this case, four hybrid orbitals are formed, oriented in space symmetrically to the four vertices of the tetrahedron, that is, they are located at an angle of 109 0 28 ". This spatial position is called tetrahedral. This structure is known for molecules NH 3, H 2 O and in general for elements of the II period. Schematically its appearance in space can be displayed in the following figure

Rice. 10. Spatial arrangement of bonds in the ammonia molecule,

projected onto a plane.

The formation of tetrahedral bonds due to sp 3 hybridization can be represented as follows (Fig. 11):

Rice. 11. Formation of tetrahedral bonds during sp 3 hybridization.

The formation of tetrahedral bonds during sp 3 hybridization using the example of a CCl 4 molecule is shown in Fig. 12.

Fig. 12. Formation of tetrahedral bonds during sp 3 - hybridization into CCl 4 molecules

Hybridization does not only concern s- and p-orbitals. To explain the stereochemical elements of III and subsequent periods, there is a need to construct hybrid orbitals simultaneously including s-, p-, d- orbitals.

Substances with covalent bonds include:

1. organic compounds;

2. solid and liquid substances in which bonds are formed between pairs of halogen atoms, as well as between pairs of hydrogen, nitrogen and oxygen atoms, for example, H2;

3. elements of group VI (for example, spiral chains of tellurium), elements of group V (for example, arsenic), elements of group IV (diamond, silicon, germanium);

4. compounds that obey the 8-N rule (such as InSb, CdS, GaAs, CdTe), when their constituent elements are located in II-VI, III-V groups in the periodic table.

IN solids With a covalent bond, different crystal structures can be formed for the same substance, the binding energy of which is almost the same. For example, the structure of ZnS can be cubic (zincblende) or hexagonal (wurtzite). The arrangement of the nearest neighbors in zinc blende and wurtzite is the same, and the only and small difference in the energies of these two structures is determined by the arrangement of the atoms next to the nearest ones. This ability of some substances is called allotropy or polymorphism. Another example of allotropy is silicon carbide, which has a number of polytypes of different structures from purely cubic to hexagonal. These numerous crystalline modifications of ZnS, SiC exist at room temperature.

Ionic bond

Ionic bonding is an electrostatic force of attraction between ions with charges of opposite sign (i.e. + and −).

The idea of ​​ionic bonding was formed on the basis of the ideas of V. Kossel. He suggested (1916) that when two atoms interact, one gives up and the other accepts electrons. Thus, an ionic bond is formed by the transfer of one or more electrons from one atom to another. For example, in sodium chloride, an ionic bond is formed by the transfer of an electron from a sodium atom to a chlorine atom. As a result of this transfer, a sodium ion with a charge of +1 and a chloride ion with a charge of -1 are formed. They are attracted to each other by electrostatic forces, forming a stable molecule. The electron transfer model proposed by Kossel allows one to explain the formation of such compounds as lithium fluoride, calcium oxide, and lithium oxide.

The most typical ionic compounds consist of metal cations belonging to groups I and II of the periodic system, and anions of non-metallic elements belonging to groups VI and VII.

The ease of formation of an ionic compound depends on the ease of formation of its constituent cations and anions. The ease of formation is higher, the lower the ionization energy of the atom donating electrons (electron donor), and the atom adding electrons (electron acceptor) has a higher affinity for the electron. Electron affinity is a measure of the ability of an atom to gain an electron. It is quantified as the change in energy that occurs when one mole of singly charged anions is formed from one mole of atoms. This is the so-called “first electron affinity” concept. The second electron affinity is the energy change that occurs when one mole of doubly charged anions is formed from one mole of singly charged anions. These concepts, that is, ionization energy and electron affinity, relate to gaseous substances and are characteristics of atoms and ions in the gaseous state. But it should be borne in mind that most ionic compounds are most stable in the solid state. This circumstance is explained by the existence of a crystal lattice in them in the solid state. The question arises. Why, after all, are ionic compounds more stable in the form of crystal lattices, and not in the gaseous state? The answer to this question is the calculation of the energy of the crystal lattice, based on the electrostatic model. In addition to this, this calculation is also a test of the theory of ionic bonding.

To calculate the energy of a crystal lattice, it is necessary to determine the work that needs to be spent on destroying the crystal lattice with the formation of gaseous ions. To carry out the calculation, the idea of ​​the forces of attraction and repulsion is used. The expression for the potential energy of interaction of singly charged ions is obtained by summing the energy of attraction and energy of repulsion

E = E in + E out (1).

The energy of Coulomb attraction of ions of opposite signs is taken as Eat, for example, Na + and Cl - for the NaCl compound

E incoming = -e 2 /4πε 0 r (2),

since the distribution of electronic charge in a filled electron shell is spherically symmetrical. Due to the repulsion that occurs due to the Pauli principle when the filled shells of the anion and cation overlap, the distance to which the ions can approach is limited. The repulsive energy changes rapidly with internuclear distance and can be written as the following two approximate expressions:

E ott = A/r n (n≈12) (3)

E ott = B∙exp(-r/ρ) (4),

where A and B are constants, r is the distance between ions, ρ is a parameter (characteristic length).

It should be noted that none of these expressions correspond to the complex quantum mechanical process that leads to repulsion.

Despite the approximate nature of these formulas, they make it possible to quite accurately calculate and accordingly describe the chemical bond in the molecules of such ionic compounds as NaCl, KCl, CaO.

Because electric field Since the ion has spherical symmetry (Fig. 13), the ionic bond, unlike the covalent bond, has no directionality. The interaction of two oppositely charged ions is compensated by repulsive forces only in the direction connecting the centers of the ion nuclei; in other directions, compensation of the electric fields of the ions does not occur. Therefore, they are able to interact with other ions. Thus, the ionic bond is not saturable.

Rice. 13. Spherical symmetry of the electrostatic field

oppositely charged charges.

Due to the non-directionality and unsaturation of ionic bonds, it is energetically most favorable when each ion is surrounded by maximum number ions of the opposite sign. Due to this, the most preferred form of existence of an ionic compound is a crystal. For example, in a NaCl crystal, each cation has six anions as its nearest neighbors.

Only at high temperatures in the gaseous state do ionic compounds exist in the form of unassociated molecules.

In ionic compounds, the coordination number does not depend on the specific electronic structure of the atoms, as in covalent compounds, but is determined by the ratio of the sizes of the ions. With the ratio ionic radii within the range of 0.41 - 0.73, octahedral coordination of ions is observed, with a ratio of 0.73-1.37 - cubic coordination, etc.

Thus, under normal conditions, ionic compounds are crystalline substances. The concept of two-ionic molecules, for example, NaCL, CsCl, does not apply to them. Each crystal consists of a large number of ions.

An ionic bond can be represented as a limiting polar bond, for which the effective charge of the atom is close to unity. For a purely covalent nonpolar bond, the effective charge of the atoms is zero. In real substances, purely ionic and purely covalent bonds are rare. Most compounds have a bond character intermediate between nonpolar covalent and polar ionic. That is, in these compounds the covalent bond is partially ionic in nature. The nature of ionic and covalent bonds in real substances is presented in Figure 14.

Rice. 14. Ionic and covalent nature of the bond.

The proportion of ionic character of a bond is called the degree of ionicity. It is characterized by the effective charges of atoms in a molecule. The degree of ionicity increases with increasing difference in electronegativity of the atoms forming it.

Metal connection

In metal atoms, the outer valence electrons are held much weaker than in non-metal atoms. This causes the loss of connection between electrons and individual atoms for a sufficiently long period of time and their socialization. A socialized ensemble of external electrons is formed. The existence of such an electronic system leads to the emergence of forces that keep positive metal ions in a close state, despite their charge of the same name. This bond is called metallic. Such a bond is characteristic only of metal and exists in the solid and liquid states of the substance. A metal bond is a type of chemical bond. It is based on the socialization of external electrons, which lose their connection with the atom and are therefore called free electrons (Fig. 15).

Rice. 15. Metal connection.

Proof of existence metal connection are the following facts. All metals have high thermal conductivity and high electrical conductivity, which is ensured by the presence of free electrons. In addition, the same circumstance determines the good reflectivity of metals to light irradiation, their shine and opacity, high ductility, positive temperature coefficient electrical resistance.

The stability of the crystal lattice of metals cannot be explained by such types of bonds as ionic and covalent. Ionic bonding between metal atoms located at the sites of the crystal lattice is impossible, since they have the same charge. Covalent bonding between metal atoms is also unlikely, since each atom has 8 to 12 nearest neighbors, and the formation of covalent bonds with so many shared electron pairs is unknown.

Metal structures are characterized by the fact that they have a rather rare arrangement of atoms (large internuclear distances) and big number the nearest neighbors of each atom in a crystal lattice. Table 1 shows three typical metal structures.

Table 1

Characteristics of the structures of the three most common metals

We see that each atom participates in the formation of a large number of bonds (for example, with 8 atoms). Such a large number of bonds (with 8 or 12 atoms) cannot be simultaneously localized in space. The connection must be carried out due to the resonance of the vibrational motion of the external electrons of each atom, as a result of which the collectivization of all external electrons of the crystal occurs with the formation of an electron gas. In many metals, to form a metallic bond, it is enough to take one electron from each atom. This is exactly what is observed for lithium, which has only one electron in its outer shell. A lithium crystal is a lattice of Li + ions (spheres with a radius of 0.068 nm) surrounded by electron gas.

Rice. 16. Various types crystal packing: a-hexagonal close packing; b - face-centered cubic packing; c-body-centered cubic packing.

There are similarities between metallic and covalent bonds. It lies in the fact that both types of bonds are based on the sharing of valence electrons. However, a covalent bond only connects two adjacent atoms, and the shared electrons are in close proximity to the bonded atoms. In a metallic bond, several atoms participate in sharing valence electrons.

Thus, the concept of a metallic bond is inextricably linked with the idea of ​​metals as a collection of positively charged ionic cores with large gaps between the ions filled with electron gas, while at the macroscopic level the system remains electrically neutral.

In addition to the types of chemical bonds discussed above, there are other types of bonds that are intermolecular: hydrogen bond, van der Waals interaction, donor-acceptor interaction.

Donor-acceptor interaction of molecules

The mechanism of formation of a covalent bond due to the two-electron cloud of one atom and the free orbital of another is called donor-acceptor. An atom or particle that provides a two-electron cloud for communication is called a donor. An atom or particle with a free orbital that accepts this electron pair is called an acceptor.

Main types of intermolecular interactions. Hydrogen bond

Between valence-saturated molecules, at distances exceeding the particle size, electrostatic forces of intermolecular attraction can appear. They are called van der Waals forces. The van der Waals interaction always exists between closely spaced atoms, but plays an important role only in the absence of stronger bonding mechanisms. This weak interaction with a characteristic energy of 0.2 eV/atom occurs between neutral atoms and between molecules. The name of the interaction is associated with the name of van der Waals, since it was he who first suggested that the equation of state, taking into account the weak interaction between gas molecules, describes the properties of real gases much better than the equation of state of an ideal gas. However, the nature of this attractive force was explained only in 1930 by London. Currently, the following three types of interactions are classified as van der Waals attraction: orientational, inductive, and dispersive (London effect). The energy of van der Waals attraction is determined by the sum of orientational, inductive and dispersion interactions.

E incoming = E or + E ind + E disp (5).

Orientation interaction (or dipole-dipole interaction) occurs between polar molecules, which, when approaching, turn (orient) towards each other with opposite poles so that the potential energy of the system of molecules becomes minimal. The greater the dipole moment of molecules μ and the smaller the distance l between them, the more significant the energy of orientation interaction is:

E or = -(μ 1 μ 2) 2 / (8π 2 ∙ε 0 ∙l 6) (6),

where ε 0 is the electrical constant.

Inductive interaction is associated with the processes of polarization of molecules by surrounding dipoles. It is more significant, the higher the polarizability α of a non-polar molecule and the greater the dipole moment μ of a polar molecule

E ind = -(αμ 2)/ (8π 2 ∙ε 0 ∙l 6) (7).

The polarizability α of a nonpolar molecule is called deformational, since it is associated with the deformation of the particle, while μ characterizes the displacement of the electron cloud and nuclei relative to their previous positions.

Dispersion interaction (London effect) occurs in any molecules, regardless of their structure and polarity. Due to the instantaneous mismatch of the centers of gravity of the charges of the electron cloud and nuclei, an instantaneous dipole is formed, which induces instantaneous dipoles in other particles. The movement of instantaneous dipoles becomes consistent. As a result, neighboring particles experience mutual attraction. The energy of dispersion interaction depends on the ionization energy E I and the polarizability of molecules α

E disp = - (E I 1 ∙E I 2)∙ α 1 α 2 /(E I 1 +E I 2) l 6 (8).

The hydrogen bond is intermediate between valence and intermolecular interactions. The hydrogen bond energy is low, 8–80 kJ/mol, but higher than the van der Waals interaction energy. Hydrogen bonding is characteristic of liquids such as water, alcohols, and acids and is caused by a positively polarized hydrogen atom. Small sizes and the absence of internal electrons allow a hydrogen atom present in a liquid in any compound to enter into additional interaction with a negatively polarized atom of another or the same molecule that is not covalently bonded to it

A δ- - H δ+…. A δ- - H δ+.

That is, an association of molecules occurs. The association of molecules leads to a decrease in volatility, an increase in the boiling point and heat of evaporation, and an increase in the viscosity and dielectric constant of liquids.

Water is a particularly suitable substance for hydrogen bonding because its molecule has two hydrogen atoms and two lone pairs on the oxygen atom. This determines the high dipole moment of the molecule (μ D = 1.86 D) and the ability to form four hydrogen bonds: two as a proton donor and two as a proton acceptor

(H 2 O….N – O…H 2 O) 2 times.

It is known from experiments that with a change in molecular weight in the series of hydrogen compounds of elements of the third and subsequent periods, the boiling point increases. If this pattern is applied to water, then its boiling point should not be 100 0 C, but 280 0 C. This contradiction confirms the existence of a hydrogen bond in water.

Experiments have shown that molecular associates are formed in liquid and especially in solid water. Ice has a tetrahedral crystal lattice. In the center of the tetrahedron there is an oxygen atom of one water molecule; at the four vertices there are oxygen atoms of neighboring molecules, which are connected by hydrogen bonds to their nearest neighbors. In liquid water, hydrogen bonds are partially destroyed, and in its structure there is a dynamic equilibrium between molecular associates and free molecules.

Valence bond method

The theory of valence bonds, or localized electron pairs, posits that each pair of atoms in a molecule is held together by one or more shared electron pairs. In the valence bond theory, a chemical bond is localized between two atoms, that is, it is two-center and two-electron.

The valence bond method is based on the following basic principles:

Each pair of atoms in a molecule is held together by one or more shared electron pairs;

A single covalent bond is formed by two electrons with antiparallel spins located on the valence orbitals of the bonding atoms;

When a bond is formed, the wave functions of electrons overlap, leading to an increase in the electron density between atoms and a decrease in the total energy of the system;

“Chemical bond” is the energy of destruction of the lattice into ions _Ekul = Uresh. Basic principles of the MO method. Types of overlap of atomic AOs. bonding and antibonding MOs with a combination of atomic orbitals s and s pz and pz px and px. H?C? C?H. ? - Repulsion coefficient. Qeff =. Ao. Basic theories of chemical bonding.

“Types of chemical bonds” - Substances with ionic bonds form an ionic crystal lattice. Atoms. Electronegativity. Municipal Educational Institution Lyceum No. 18 chemistry teacher Kalinina L.A. Ions. For example: Na1+ and Cl1-, Li1+ and F1- Na1+ + Cl1- = Na(:Cl:) . If e - are added, the ion becomes negatively charged. The atomic frame has high strength.

“The Life of Mendeleev” - July 18 D.I. Mendeleev graduated from the Tobolsk gymnasium. August 9, 1850 - June 20, 1855 while studying in Main Pedagogical Institute. “If you do not know names, then the knowledge of things will die” K. Liney. Life and work of D.I. Mendeleev. Ivan Pavlovich Mendeleev (1783 - 1847), father of the scientist. Discovery of the periodic law.

“Types of chemical bonds” - H3N. Al2O3. The structure of matter." H2S. MgO. H2. Cu. Mg S.CS2. I. Write down the formulas of the substances: 1.c.N.S. 2.s K.P.S. 3. with I.S. K.N.S. NaF. C.K.P.S. Determine the type of chemical bond. Which of the molecules corresponds to the scheme: A A?

"Mendeleev" - Dobereiner's Triads of Elements. Gases. Work. Life and scientific feat. Periodic table of elements (long form). Newlands' "Law of Octaves" Scientific activity. Solutions. A new stage of life. The second version of Mendeleev's system of elements. Part of L. Meyer's table of elements. Discovery of the periodic law (1869).

“The Life and Work of Mendeleev” - Ivan Pavlovich Mendeleev (1783 - 1847), the scientist’s father. 1834, January 27 (February 6) - D.I. Mendeleev was born in the city of Tobolsk, in Siberia. 1907, January 20 (February 2) D.I. Mendeleev died of heart paralysis. DI. Menedeleev (South Kazakhstan region, Shymkent city). Industry. On July 18, 1849, D.I. Mendeleev graduated from the Tobolsk gymnasium.

Characteristics of chemical bonds

The doctrine of chemical bonding forms the basis of all theoretical chemistry. A chemical bond is understood as the interaction of atoms that binds them into molecules, ions, radicals, and crystals. There are four types of chemical bonds: ionic, covalent, metallic and hydrogen. Different types of bonds can be found in the same substances.

1. In bases: between the oxygen and hydrogen atoms in hydroxo groups the bond is polar covalent, and between the metal and the hydroxo group it is ionic.

2. In salts of oxygen-containing acids: between the non-metal atom and the oxygen of the acidic residue - covalent polar, and between the metal and the acidic residue - ionic.

3. In ammonium, methylammonium, etc. salts, between the nitrogen and hydrogen atoms there is a polar covalent, and between ammonium or methylammonium ions and the acid residue - ionic.

4. In metal peroxides (for example, Na 2 O 2), the bond between the oxygen atoms is covalent, nonpolar, and between the metal and oxygen is ionic, etc.

The reason for the unity of all types and types of chemical bonds is their identical chemical nature - electron-nuclear interaction. The formation of a chemical bond in any case is the result of electron-nuclear interaction of atoms, accompanied by the release of energy.

Methods for forming a covalent bond

Covalent chemical bond is a bond that arises between atoms due to the formation of shared electron pairs.

Covalent compounds are usually gases, liquids, or relatively low-melting solids. One of the rare exceptions is diamond, which melts above 3,500 °C. This is explained by the structure of diamond, which is a continuous lattice of covalently bonded carbon atoms, and not a collection of individual molecules. In fact, any diamond crystal, regardless of its size, is one huge molecule.

A covalent bond occurs when the electrons of two nonmetal atoms combine. The resulting structure is called a molecule.

The mechanism of formation of such a bond can be exchange or donor-acceptor.

In most cases, two covalently bonded atoms have different electronegativity and the shared electrons do not belong to the two atoms equally. Most of the time they are closer to one atom than to another. In a hydrogen chloride molecule, for example, the electrons that form a covalent bond are located closer to the chlorine atom because its electronegativity is higher than that of hydrogen. However, the difference in the ability to attract electrons is not large enough for complete electron transfer from the hydrogen atom to the chlorine atom to occur. Therefore, the bond between hydrogen and chlorine atoms can be considered as a cross between an ionic bond (complete electron transfer) and a non-polar covalent bond (a symmetrical arrangement of a pair of electrons between two atoms). The partial charge on atoms is denoted by the Greek letter δ. Such a bond is called a polar covalent bond, and the hydrogen chloride molecule is said to be polar, that is, it has a positively charged end (hydrogen atom) and a negatively charged end (chlorine atom).

1. The exchange mechanism operates when atoms form shared electron pairs by combining unpaired electrons.

1) H 2 - hydrogen.

The bond occurs due to the formation of a common electron pair by the s-electrons of hydrogen atoms (overlapping s-orbitals).

2) HCl - hydrogen chloride.

The bond occurs due to the formation of a common electron pair of s- and p-electrons (overlapping s-p orbitals).

3) Cl 2: In a chlorine molecule, a covalent bond is formed due to unpaired p-electrons (overlapping p-p orbitals).

4) N ​​2: In the nitrogen molecule, three common electron pairs are formed between the atoms.

Donor-acceptor mechanism of covalent bond formation

Donor has an electron pair acceptor- free orbital that this pair can occupy. In the ammonium ion, all four bonds with hydrogen atoms are covalent: three were formed due to the creation of common electron pairs by the nitrogen atom and hydrogen atoms according to the exchange mechanism, one - through the donor-acceptor mechanism. Covalent bonds are classified by the way the electron orbitals overlap, as well as by their displacement towards one of the bonded atoms. Chemical bonds formed as a result of overlapping electron orbitals along a bond line are called σ - connections(sigma bonds). The sigma bond is very strong.

The p orbitals can overlap in two regions, forming a covalent bond through lateral overlap.

Chemical bonds formed as a result of the “lateral” overlap of electron orbitals outside the bond line, i.e., in two regions, are called pi bonds.

According to the degree of displacement of common electron pairs to one of the atoms they connect, a covalent bond can be polar or non-polar. A covalent chemical bond formed between atoms with the same electronegativity is called non-polar. Electron pairs are not displaced towards any of the atoms, since atoms have the same electronegativity - the property of attracting valence electrons from other atoms. For example,

i.e., molecules are formed through a covalent nonpolar bond simple substances-non-metals. A covalent chemical bond between atoms of elements whose electronegativity differs is called polar.

For example, NH 3 is ammonia. Nitrogen is a more electronegative element than hydrogen, so the shared electron pairs are shifted towards its atom.

Characteristics of a covalent bond: bond length and energy

The characteristic properties of a covalent bond are its length and energy. Bond length is the distance between atomic nuclei. The shorter the length of a chemical bond, the stronger it is. However, a measure of bond strength is bond energy, which is determined by the amount of energy required to break the bond. It is usually measured in kJ/mol. Thus, according to experimental data, the bond lengths of the H 2, Cl 2 and N 2 molecules, respectively, are 0.074, 0.198 and 0.109 nm, and the bond energies, respectively, are 436, 242 and 946 kJ/mol.

Ions. Ionic bond

There are two main possibilities for an atom to obey the octet rule. The first of these is the formation of ionic bonds. (The second is the formation of a covalent bond, which will be discussed below). When an ionic bond is formed, a metal atom loses electrons, and a non-metal atom gains electrons.

Let's imagine that two atoms “meet”: an atom of a group I metal and a non-metal atom of group VII. A metal atom has a single electron at its outer energy level, while a non-metal atom just lacks one electron for its outer level to be complete. The first atom will easily give the second its electron, which is far from the nucleus and weakly bound to it, and the second will provide it with a free place on its outer electronic level. Then the atom, deprived of one of its negative charges, will become a positively charged particle, and the second will turn into a negatively charged particle due to the resulting electron. Such particles are called ions.

This is a chemical bond that occurs between ions. Numbers showing the number of atoms or molecules are called coefficients, and numbers showing the number of atoms or ions in a molecule are called indices.

Metal connection

Metals have specific properties, different from the properties of other substances. Such properties are relatively high melting temperatures, the ability to reflect light, and high thermal and electrical conductivity. These features are due to the existence in metals special type connection - metal connection.

Metallic bond is a bond between positive ions in metal crystals, carried out due to the attraction of electrons moving freely throughout the crystal. The atoms of most metals at the outer level contain a small number of electrons - 1, 2, 3. These electrons come off easily, and the atoms turn into positive ions. The detached electrons move from one ion to another, binding them into a single whole. Connecting with ions, these electrons temporarily form atoms, then break off again and combine with another ion, etc. A process occurs endlessly, which can be schematically depicted as follows:

Consequently, in the volume of the metal, atoms are continuously converted into ions and vice versa. The bond in metals between ions through shared electrons is called metallic. The metallic bond has some similarities with the covalent bond, since it is based on the sharing of external electrons. However, with a covalent bond, the outer unpaired electrons of only two neighboring atoms are shared, while with a metallic bond, all atoms take part in the sharing of these electrons. That is why crystals with a covalent bond are brittle, but with a metal bond, as a rule, they are ductile, electrically conductive and have a metallic luster.

Metallic bonding is characteristic of both pure metals and mixtures of various metals - alloys in solid and liquid states. However, in the vapor state, metal atoms are connected to each other by a covalent bond (for example, sodium vapor fills yellow light lamps to illuminate the streets of large cities). Metal pairs consist of individual molecules (monatomic and diatomic).

A metal bond also differs from a covalent bond in strength: its energy is 3-4 times less than the energy of a covalent bond.

Bond energy is the energy required to break a chemical bond in all molecules that make up one mole of a substance. The energies of covalent and ionic bonds are usually high and amount to values ​​of the order of 100-800 kJ/mol.

Hydrogen bond

Chemical bond between positively polarized hydrogen atoms of one molecule(or parts thereof) and negatively polarized atoms of highly electronegative elements having shared electron pairs (F, O, N and less often S and Cl), another molecule (or parts thereof) is called hydrogen. The mechanism of hydrogen bond formation is partly electrostatic, partly d honoror-acceptor character.

Examples of intermolecular hydrogen bonding:

In the presence of such a connection, even low-molecular substances can, under normal conditions, be liquids (alcohol, water) or easily liquefied gases (ammonia, hydrogen fluoride). In biopolymers - proteins (secondary structure) - there is an intramolecular hydrogen bond between carbonyl oxygen and the hydrogen of the amino group:

Polynucleotide molecules - DNA (deoxyribonucleic acid) - are double helices in which two chains of nucleotides are linked to each other by hydrogen bonds. In this case, the principle of complementarity operates, i.e., these bonds are formed between certain pairs consisting of purine and pyrimidine bases: the thymine (T) is located opposite the adenine nucleotide (A), and the cytosine (C) is located opposite the guanine (G).

Substances with hydrogen bonds have molecular crystal lattices.

There is no unified theory of chemical bonds; chemical bonds are conventionally divided into covalent (a universal type of bond), ionic (a special case of a covalent bond), metallic and hydrogen.

Covalent bond

The formation of a covalent bond is possible by three mechanisms: exchange, donor-acceptor and dative (Lewis).

According to metabolic mechanism The formation of a covalent bond occurs due to the sharing of common electron pairs. In this case, each atom tends to acquire a shell of an inert gas, i.e. obtain a completed external energy level. The formation of a chemical bond by exchange type is depicted using Lewis formulas, in which each valence electron of an atom is represented by dots (Fig. 1).

Rice. 1 Formation of a covalent bond in the HCl molecule by the exchange mechanism

With the development of the theory of atomic structure and quantum mechanics, the formation of a covalent bond is represented as the overlap of electronic orbitals (Fig. 2).

Rice. 2. Formation of a covalent bond due to the overlap of electron clouds

The greater the overlap of atomic orbitals, the stronger the bond, the shorter the bond length, and the greater the bond energy. A covalent bond can be formed by overlapping different orbitals. As a result of the overlap of s-s, s-p orbitals, as well as d-d, p-p, d-p orbitals with lateral lobes, the formation of bonds occurs. A bond is formed perpendicular to the line connecting the nuclei of 2 atoms. One - and one - bond are capable of forming a multiple (double) covalent bond, characteristic of organic matter class of alkenes, alkadienes, etc. One and two bonds form a multiple (triple) covalent bond, characteristic of organic substances of the class of alkynes (acetylenes).

Formation of a covalent bond by donor-acceptor mechanism Let's look at the example of the ammonium cation:

NH 3 + H + = NH 4 +

7 N 1s 2 2s 2 2p 3

The nitrogen atom has a free lone pair of electrons (electrons not involved in the formation of chemical bonds within the molecule), and the hydrogen cation has a free orbital, so they are an electron donor and acceptor, respectively.

Let us consider the dative mechanism of covalent bond formation using the example of a chlorine molecule.

17 Cl 1s 2 2s 2 2p 6 3s 2 3p 5

The chlorine atom has both a free lone pair of electrons and vacant orbitals, therefore, it can exhibit the properties of both a donor and an acceptor. Therefore, when a chlorine molecule is formed, one chlorine atom acts as a donor and the other as an acceptor.

Main characteristics of a covalent bond are: saturation (saturated bonds are formed when an atom attaches as many electrons to itself as its valence capabilities allow; unsaturated bonds are formed when the number of attached electrons is less than the valence capabilities of the atom); directionality (this value is related to the geometry of the molecule and the concept of “bond angle” - the angle between bonds).

Ionic bond

There are no compounds with a pure ionic bond, although this is understood as a chemically bonded state of atoms in which a stable electronic environment of the atom is created when the total electron density is completely transferred to the atom of a more electronegative element. Ionic bonding is possible only between atoms of electronegative and electropositive elements that are in the state of oppositely charged ions - cations and anions.

DEFINITION

Ion are electrically charged particles formed by the removal or addition of an electron to an atom.

When transferring an electron, metal and nonmetal atoms tend to form a stable electron shell configuration around their nucleus. A non-metal atom creates a shell of the subsequent inert gas around its core, and a metal atom creates a shell of the previous inert gas (Fig. 3).

Rice. 3. Formation of an ionic bond using the example of a sodium chloride molecule

Molecules in which ionic bonds exist in their pure form are found in the vapor state of the substance. The ionic bond is very strong, and therefore substances with this bond have a high melting point. Unlike covalent bonds, ionic bonds are not characterized by directionality and saturation, since the electric field created by ions acts equally on all ions due to spherical symmetry.

Metal connection

The metallic bond is realized only in metals - this is the interaction that holds metal atoms in a single lattice. Only the valence electrons of the metal atoms belonging to its entire volume participate in the formation of a bond. In metals, electrons are constantly stripped from atoms and move throughout the entire mass of the metal. Metal atoms, deprived of electrons, turn into positively charged ions, which tend to accept moving electrons. This continuous process forms the so-called “electron gas” inside the metal, which firmly binds all the metal atoms together (Fig. 4).

The metallic bond is strong, so metals are characterized heat melting, and the presence of “electron gas” gives metals malleability and ductility.

Hydrogen bond

A hydrogen bond is a specific intermolecular interaction, because its occurrence and strength depend on the chemical nature of the substance. It is formed between molecules in which a hydrogen atom is bonded to an atom with high electronegativity (O, N, S). The occurrence of a hydrogen bond depends on two reasons: firstly, the hydrogen atom associated with an electronegative atom does not have electrons and can easily be incorporated into the electron clouds of other atoms, and, secondly, having a valence s-orbital, the hydrogen atom is able to accept a lone pair electrons of an electronegative atom and form a bond with it through the donor-acceptor mechanism.

1. Alkaline earth metals are

5) to s-elements

6) to p-elements

7) to d-elements

8) to f - elements

2. How many electrons do atoms of alkaline earth metals contain at the outer energy level?

1) One 2) two 3) three 4) four

3. B chemical reactions aluminum atoms exhibit

3) Oxidizing properties 2) acidic properties

4) 3) restorative properties 4) basic properties

4. The interaction of calcium with chlorine is a reaction

1) Decomposition 2) connection 3) substitution 4) exchange

5. The molecular weight of sodium bicarbonate is:

1) 84 2) 87 3) 85 4) 86

3. Which atom is heavier - iron or silicon - and by how much?

4. Determine the relative molecular weights of simple substances: hydrogen, oxygen, chlorine, copper, diamond (carbon). Remember which of them consist of diatomic molecules, and which of atoms.
5.calculate the relative molecular masses of the following compounds: carbon dioxide CO2 sulfuric acid H2SO4 sugar C12H22O11 ethyl alcohol C2H6O marble CaCPO3
6.In hydrogen peroxide, there is one hydrogen atom for every oxygen atom. Determine the formula of hydrogen preoxide if it is known that its relative molecular weight is 34. What is the mass ratio of hydrogen and oxygen in this compound?
7. How many times is a carbon dioxide molecule heavier than an oxygen molecule?

Please help me, 8th grade assignment.