Covalent bond formed by the interaction of nonmetals. Nonmetal atoms have high electronegativity and tend to fill the outer electron layer with foreign electrons. Two such atoms can go into a stable state if they combine their electrons .

Let us consider the formation of a covalent bond in simple substances.

1.Formation of a hydrogen molecule.

Every atom hydrogen has one electron. To transition to a stable state, it needs one more electron.

When two atoms come close, the electron clouds overlap. A shared electron pair is formed, which bonds the hydrogen atoms into a molecule.

The space between two nuclei shares more electrons than other places. An area with increased electron density and negative charge. Positively charged nuclei are attracted to it, and a molecule is formed.

In this case, each atom receives a completed two-electron outer level and goes into a stable state.

A covalent bond due to the formation of one shared electron pair is called single.

Shared electron pairs (covalent bonds) are formed due to unpaired electrons, located on the outer energy levels of interacting atoms.

Hydrogen has one unpaired electron. For other elements, their number is 8 - group number.

Nonmetals VII And groups (halogens) have one unpaired electron on the outer layer.

In non-metals VI A groups (oxygen, sulfur) have two such electrons.

In non-metals V And groups (nitrogen, phosphorus) have three unpaired electrons.

2.Formation of a fluorine molecule.

Atom fluoride has seven electrons in the outer level. Six of them form pairs, and the seventh is unpaired.

When atoms join, one common electron pair is formed, that is, one covalent bond occurs. Each atom receives a completed eight-electron outer layer. The bond in the fluorine molecule is also single. The same single bonds exist in molecules chlorine, bromine and iodine .

If atoms have several unpaired electrons, then two or three common pairs are formed.

3.Formation of an oxygen molecule.

At the atom oxygen at the outer level there are two unpaired electrons.

When two atoms interact oxygen two common electron pairs arise. Each atom fills its outer level with up to eight electrons. The oxygen molecule has a double bond.

A covalent bond occurs due to the sharing of electrons belonging to both atoms participating in the interaction. The electronegativity of nonmetals is large enough that no electron transfer occurs.

Electrons in overlapping electron orbitals are shared. This creates a situation in which the outer electronic levels of the atoms are filled, that is, an 8- or 2-electron outer shell is formed.

The state in which the electron shell is completely filled is characterized by the lowest energy and, accordingly, maximum stability.

There are two mechanisms of formation:

1. donor-acceptor;
2. exchange.

In the first case, one of the atoms provides its pair of electrons, and the second provides a free electron orbital.

In the second, one electron from each participant in the interaction comes into the common pair.

Depending on what type they are- atomic or molecular, compounds with a similar type of bond can vary significantly in physicochemical characteristics.

Molecular substances most often gases, liquids or solids with low melting and boiling points, non-electrically conductive, and low strength. These include: hydrogen (H 2), oxygen (O 2), nitrogen (N 2), chlorine (Cl 2), bromine (Br 2), orthorhombic sulfur (S 8), white phosphorus (P 4) and others simple substances; carbon dioxide (CO 2), sulfur dioxide (SO 2), nitrogen oxide V (N 2 O 5), water (H 2 O), hydrogen chloride (HCl), hydrogen fluoride (HF), ammonia (NH 3), methane (CH 4), ethyl alcohol (C 2 H 5 OH), organic polymers and others.

Atomic substances exist in the form of durable crystals with high boiling and melting points, are insoluble in water and other solvents, and many do not conduct electric current. An example is diamond, which has exceptional strength. This is explained by the fact that diamond is a crystal consisting of carbon atoms connected by covalent bonds. There are no individual molecules in a diamond. Also, substances such as graphite, silicon (Si), silicon dioxide (SiO 2), silicon carbide (SiC) and others have an atomic structure.

Covalent bonds can be not only single (as in the chlorine molecule Cl2), but also double, as in the oxygen molecule O2, or triple, as, for example, in the nitrogen molecule N2. At the same time, triple ones have more energy and are more durable than double and single ones.

A covalent bond can be formed both between two atoms of the same element (non-polar) and between atoms of different chemical elements (polar).

It is not difficult to indicate the formula of a compound with a covalent polar bond if you compare the electronegativity values ​​of the atoms that make up the molecules. No difference in electronegativity will determine non-polarity. If there is a difference, then the molecule will be polar.

Don't miss: mechanism of education, specific examples.

Covalent nonpolar chemical bond

Characteristic of simple substances, non-metals. The electrons belong to the atoms equally, and there is no shift in the electron density.

Examples include the following molecules:

H2, O2, O3, N2, F2, Cl2.

The exception is inert gases. Their outer energy level is completely filled, and the formation of molecules is energetically unfavorable for them, and therefore they exist in the form of individual atoms.

Also, an example of substances with a non-polar covalent bond would be, for example, PH3. Despite the fact that the substance consists of different elements, the electronegativities of the elements do not actually differ, which means that the electron pair will not shift.

Covalent polar chemical bond

Considering a covalent polar bond, many examples can be given: HCl, H2O, H2S, NH3, CH4, CO2, SO3, CCl4, SiO2, CO.

formed between nonmetal atoms with different electronegativity. In this case, the nucleus of an element with greater electronegativity attracts shared electrons closer to itself.

Scheme of formation of a polar covalent bond

Depending on the mechanism of formation, they can become common electrons of one or both atoms.

The picture clearly shows the interaction in the hydrochloric acid molecule.

A pair of electrons belongs to both one atom and the second, both of them, thus, the outer levels are filled. But the more electronegative chlorine attracts a pair of electrons a little closer to itself (while it remains shared). The difference in electronegativity is not large enough for a pair of electrons to go completely to one of the atoms. As a result, a partial negative charge appears on chlorine and a partial positive charge on hydrogen. The HCl molecule is a polar molecule.

Physico-chemical properties of the bond

The connection can be characterized by the following properties: directivity, polarity, polarizability and saturation.

Atoms of most elements do not exist separately, as they can interact with each other. This interaction produces more complex particles.

The nature of a chemical bond is the action of electrostatic forces, which are the forces of interaction between electric charges. Electrons and atomic nuclei have such charges.

Electrons located on the outer electronic levels (valence electrons), being farthest from the nucleus, interact with it weakest, and therefore are able to break away from the nucleus. They are responsible for bonding atoms to each other.

Types of interactions in chemistry

Types of chemical bonds can be presented in the following table:

Characteristics of ionic bonding

Chemical reaction that occurs due to ion attraction having different charges is called ionic. This happens if the atoms being bonded have a significant difference in electronegativity (that is, the ability to attract electrons) and the electron pair goes to the more electronegative element. The result of this transfer of electrons from one atom to another is the formation of charged particles - ions. An attraction arises between them.

They have the lowest electronegativity indices typical metals, and the largest are typical non-metals. Ions are thus formed by the interaction between typical metals and typical nonmetals.

Metal atoms become positively charged ions (cations), donating electrons to their outer electron levels, and nonmetals accept electrons, thus turning into negatively charged ions (anions).

Atoms move into a more stable energy state, completing their electronic configurations.

The ionic bond is non-directional and non-saturable, since the electrostatic interaction occurs in all directions; accordingly, the ion can attract ions of the opposite sign in all directions.

The arrangement of the ions is such that around each there is a certain number of oppositely charged ions. The concept of "molecule" for ionic compounds doesn't make sense.

Examples of education

The formation of a bond in sodium chloride (nacl) is due to the transfer of an electron from the Na atom to the Cl atom to form the corresponding ions:

Na 0 - 1 e = Na + (cation)

Cl 0 + 1 e = Cl - (anion)

In sodium chloride, there are six chloride anions around the sodium cations, and six sodium ions around each chloride ion.

When interaction is formed between atoms in barium sulfide, the following processes occur:

Ba 0 - 2 e = Ba 2+

S 0 + 2 e = S 2-

Ba donates its two electrons to sulfur, resulting in the formation of sulfur anions S 2- and barium cations Ba 2+.

Metal chemical bond

The number of electrons in the outer energy levels of metals is small; they are easily separated from the nucleus. As a result of this detachment, metal ions and free electrons are formed. These electrons are called "electron gas". Electrons move freely throughout the volume of the metal and are constantly bound and separated from atoms.

The structure of the metal substance is as follows: the crystal lattice is the skeleton of the substance, and between its nodes electrons can move freely.

The following examples can be given:

Mg - 2е<->Mg 2+

Cs-e<->Cs+

Ca - 2e<->Ca2+

Fe-3e<->Fe 3+

Covalent: polar and non-polar

The most common type of chemical interaction is a covalent bond. The electronegativity values ​​of the elements that interact do not differ sharply; therefore, only a shift of the common electron pair to a more electronegative atom occurs.

Covalent interactions can be formed by an exchange mechanism or a donor-acceptor mechanism.

The exchange mechanism is realized if each of the atoms has unpaired electrons on the outer electronic levels and the overlap of atomic orbitals leads to the appearance of a pair of electrons that already belongs to both atoms. When one of the atoms has a pair of electrons on the outer electronic level, and the other has a free orbital, then when the atomic orbitals overlap, the electron pair is shared and interacts according to the donor-acceptor mechanism.

Covalent ones are divided by multiplicity into:

• simple or single;
• double;
• triples.

Double ones ensure the sharing of two pairs of electrons at once, and triple ones - three.

According to the distribution of electron density (polarity) between bonded atoms, a covalent bond is divided into:

• non-polar;
• polar.

A nonpolar bond is formed by identical atoms, and a polar bond is formed by different electronegativity.

The interaction of atoms with similar electronegativity is called a nonpolar bond. The common pair of electrons in such a molecule is not attracted to either atom, but belongs equally to both.

The interaction of elements differing in electronegativity leads to the formation of polar bonds. In this type of interaction, shared electron pairs are attracted to the more electronegative element, but are not completely transferred to it (that is, the formation of ions does not occur). As a result of this shift in electron density, partial charges appear on the atoms: the more electronegative one has a negative charge, and the less electronegative one has a positive charge.

Properties and characteristics of covalency

Main characteristics of a covalent bond:

• The length is determined by the distance between the nuclei of interacting atoms.
• Polarity is determined by the displacement of the electron cloud towards one of the atoms.
• Directionality is the property of forming bonds oriented in space and, accordingly, molecules having certain geometric shapes.
• Saturation is determined by the ability to form a limited number of bonds.
• Polarizability is determined by the ability to change polarity under the influence of an external electric field.
• The energy required to break a bond determines its strength.

An example of a covalent non-polar interaction can be the molecules of hydrogen (H2), chlorine (Cl2), oxygen (O2), nitrogen (N2) and many others.

H· + ·H → H-H molecule has a single non-polar bond,

O: + :O → O=O molecule has a double nonpolar,

Ṅ: + Ṅ: → N≡N the molecule is triple nonpolar.

Examples of covalent bonds of chemical elements include molecules of carbon dioxide (CO2) and carbon monoxide (CO), hydrogen sulfide (H2S), hydrochloric acid (HCL), water (H2O), methane (CH4), sulfur oxide (SO2) and many others .

In the CO2 molecule, the relationship between carbon and oxygen atoms is covalent polar, since the more electronegative hydrogen attracts electron density. Oxygen has two unpaired electrons in its outer shell, while carbon can provide four valence electrons to form the interaction. As a result, double bonds are formed and the molecule looks like this: O=C=O.

In order to determine the type of bond in a particular molecule, it is enough to consider its constituent atoms. Simple metal substances form a metallic bond, metals with nonmetals form an ionic bond, simple nonmetal substances form a covalent nonpolar bond, and molecules consisting of different nonmetals form through a polar covalent bond.

The formation of chemical compounds is due to the emergence of chemical bonds between atoms in molecules and crystals.

A chemical bond is the mutual adhesion of atoms in a molecule and a crystal lattice as a result of the action of electric forces of attraction between the atoms.

COVALENT BOND.

A covalent bond is formed due to shared electron pairs that appear in the shells of the bonded atoms. It can be formed by atoms of the same element and then it non-polar; for example, such a covalent bond exists in the molecules of single-element gases H2, O2, N2, Cl2, etc.

A covalent bond can be formed by atoms of different elements that are similar in chemical character, and then it polar; for example, such a covalent bond exists in the molecules H2O, NF3, CO2. A covalent bond is formed between atoms of elements,

Quantitative characteristics of chemical bonds. Energy of communication. Link length. Polarity of a chemical bond. Bond angle. Effective charges on atoms in molecules. Dipole moment of a chemical bond. Dipole moment of a polyatomic molecule. Factors that determine the magnitude of the dipole moment of a polyatomic molecule.

Characteristics of a covalent bond . Important quantitative characteristics of a covalent bond are bond energy, its length and dipole moment.

Communication energy- the energy released during its formation, or required to separate two bonded atoms. The bond energy characterizes its strength.

Link length- the distance between the centers of bonded atoms. The shorter the length, the stronger the chemical bond.

Dipole moment of coupling(m) is a vector quantity characterizing the polarity of the connection.

The length of the vector is equal to the product of the bond length l and the effective charge q, which atoms acquire when the electron density shifts: | m | = lХ q. The dipole moment vector is directed from the positive charge to the negative one. By vectorial addition of the dipole moments of all bonds, the dipole moment of the molecule is obtained.

The characteristics of bonds are affected by their multiplicity:

The binding energy increases in a series;

The length of the connection increases in the reverse order.

Communication energy(for a given state of the system) - the difference between the energy of the state in which the constituent parts of the system are infinitely distant from each other and are in a state of active rest and the total energy of the bound state of the system: ,

where E is the binding energy of components in a system of N components (particles), Ei is the total energy of the i-th component in an unbound state (an infinitely distant particle at rest) and E is the total energy of a bound system. For a system consisting of infinitely distant particles at rest, the binding energy is usually considered equal to zero, that is, when a bound state is formed, energy is released. The binding energy is equal to the minimum work that must be expended to decompose the system into its constituent particles.

It characterizes the stability of the system: the higher the binding energy, the more stable the system. For valence electrons (electrons of the outer electron shells) of neutral atoms in the ground state, the binding energy coincides with the ionization energy, for negative ions - with the electron affinity. The chemical bond energy of a diatomic molecule corresponds to the energy of its thermal dissociation, which is on the order of hundreds of kJ/mol. The binding energy of hadrons in the atomic nucleus is determined mainly by the strong interaction. For light nuclei it is ~0.8 MeV per nucleon.

Chemical bond length— the distance between the nuclei of chemically bonded atoms. The length of a chemical bond is an important physical quantity that determines the geometric dimensions of a chemical bond and its extent in space. Various methods are used to determine the length of a chemical bond. Gas electron diffraction, microwave spectroscopy, Raman spectra, and high-resolution IR spectra are used to estimate the chemical bond lengths of isolated molecules in the vapor (gas) phase. It is believed that the length of a chemical bond is an additive quantity determined by the sum of the covalent radii of the atoms that make up the chemical bond.

Polarity of chemical bonds- a characteristic of a chemical bond, showing a change in the distribution of electron density in the space around the nuclei in comparison with the distribution of electron density in the neutral atoms forming this bond. It is possible to quantify the polarity of a bond in a molecule. The difficulty of an accurate quantitative assessment is that the polarity of the bond depends on several factors: the size of the atoms and ions of the connecting molecules; from the number and nature of the connections that the connecting atoms already had before their given interaction; on the type of structure and even the characteristics of defects in their crystal lattices. These kinds of calculations are made by various methods, which, in general, give approximately the same results (values).

For example, for HCl it has been established that each of the atoms in this molecule has a charge equal to 0.17 of the charge of a whole electron. On the hydrogen atom is +0.17, and on the chlorine atom -0.17. The so-called effective charges on atoms are most often used as a quantitative measure of bond polarity. The effective charge is defined as the difference between the charge of electrons located in some region of space near the nucleus and the charge of the nucleus. However, this measure has only a conditional and approximate [relative] meaning, since it is impossible to unambiguously identify a region in a molecule that relates exclusively to an individual atom, and in the case of several bonds, to a specific bond.

Bond angle- the angle formed by the directions of chemical (covalent) bonds emanating from one atom. Knowledge of bond angles is necessary to determine the geometry of molecules. Bond angles depend both on the individual characteristics of the attached atoms and on the hybridization of the atomic orbitals of the central atom. For simple molecules, the bond angle, like other geometric parameters of the molecule, can be calculated using quantum chemistry methods. They are determined experimentally from the values ​​of the moments of inertia of molecules obtained by analyzing their rotational spectra. The bond angle of complex molecules is determined by diffraction structural analysis methods.

EFFECTIVE CHARGE OF AN ATOM, characterizes the difference between the number of electrons belonging to a given atom in a chemical. conn., and the number of free electrons. atom. For assessments of E. z. A. they use models in which experimentally determined quantities are represented as functions of point non-polarizable charges localized on atoms; for example, the dipole moment of a diatomic molecule is considered as the product of E. z. A. to the interatomic distance. Within the framework of such models, E. z. A. can be calculated using optical data. or X-ray spectroscopy.

Dipole moments of molecules.

An ideal covalent bond exists only in particles consisting of identical atoms (H2, N2, etc.). If a bond is formed between different atoms, then the electron density shifts to one of the atomic nuclei, that is, polarization of the bond occurs. The polarity of a bond is characterized by its dipole moment.

The dipole moment of a molecule is equal to the vector sum of the dipole moments of its chemical bonds. If polar bonds are arranged symmetrically in a molecule, then the positive and negative charges cancel each other out, and the molecule as a whole is nonpolar. This happens, for example, with a molecule of carbon dioxide. Polyatomic molecules with an asymmetrical arrangement of polar bonds are generally polar. This applies in particular to the water molecule.

The resulting dipole moment of a molecule can be affected by the lone pair of electrons. Thus, NH3 and NF3 molecules have a tetrahedral geometry (taking into account the lone pair of electrons). The degrees of ionicity of the nitrogen–hydrogen and nitrogen–fluorine bonds are 15 and 19%, respectively, and their lengths are 101 and 137 pm, respectively. Based on this, one could conclude that NF3 has a larger dipole moment. However, experiment shows the opposite. For a more accurate prediction of the dipole moment, the direction of the dipole moment of the lone pair should be taken into account (Fig. 29).

The concept of hybridization of atomic orbitals and the spatial structure of molecules and ions. Features of the electron density distribution of hybrid orbitals. Main types of hybridization: sp, sp2, sp3, dsp2, sp3d, sp3d2. Hybridization involving lone electron pairs.

HYBRIDIZATION OF ATOMIC ORBITALS.

To explain the structure of some molecules, the BC method uses the atomic orbital (AO) hybridization model. For some elements (beryllium, boron, carbon), both s- and p-electrons take part in the formation of covalent bonds. These electrons are located on AOs that differ in shape and energy. Despite this, the connections formed with their participation turn out to be of equal value and located symmetrically.

In the molecules BeC12, BC13 and CC14, for example, the bond angle C1-E-C1 is 180, 120, and 109.28 o. The values ​​and energies of the E-C1 bond lengths are the same for each of these molecules. The principle of orbital hybridization is that the original AOs of different shapes and energies, when mixed, give new orbitals of the same shape and energy. The type of hybridization of the central atom determines the geometric shape of the molecule or ion formed by it.

Let us consider the structure of the molecule from the standpoint of hybridization of atomic orbitals.

Spatial shape of molecules.

Lewis formulas say a lot about the electronic structure and stability of molecules, but so far they cannot say anything about their spatial structure. In chemical bonding theory, there are two good approaches to explaining and predicting molecular geometry. They agree well with each other. The first approach is called the theory of valence electron pair repulsion (VEP). Despite the “scary” name, the essence of this approach is very simple and clear: chemical bonds and lone electron pairs in molecules tend to be located as far from each other as possible. Let us explain with specific examples. There are two Be-Cl bonds in the BeCl2 molecule. The shape of this molecule should be such that both of these bonds and the chlorine atoms at their ends are located as far apart as possible:

This is possible only with a linear form of the molecule, when the angle between the bonds (ClBeCl angle) is 180°.

Another example: the BF3 molecule has 3 B-F bonds. They are located as far apart as possible and the molecule has the shape of a flat triangle, where all angles between bonds (FBF angles) are equal to 120 o:

Hybridization of atomic orbitals.

Hybridization involves not only bonding electrons, but also lone electron pairs . For example, a water molecule contains two covalent chemical bonds between an oxygen atom and two hydrogen atoms (Figure 21).

In addition to the two pairs of electrons shared with hydrogen atoms, the oxygen atom has two pairs of outer electrons that do not participate in bond formation ( lone electron pairs). All four pairs of electrons occupy specific regions in the space around the oxygen atom. Because electrons repel each other, electron clouds are located as far apart as possible. In this case, as a result of hybridization, the shape of the atomic orbitals changes; they are elongated and directed towards the vertices of the tetrahedron. Therefore, the water molecule has an angular shape, and the angle between the oxygen-hydrogen bonds is 104.5 o.

The shape of molecules and ions of type AB2, AB3, AB4, AB5, AB6. d-AOs involved in the formation of σ bonds in flat square molecules, in octahedral molecules and in molecules built in the form of a trigonal bipyramid. The influence of repulsion of electron pairs on the spatial configuration of molecules (the concept of the participation of lone electron pairs of KNEP).

Form of molecules and ions of type AB2, AB3, AB4, AB5, AB6. Each type of AO hybridization corresponds to a strictly defined geometric shape, confirmed experimentally. Its basis is created by σ-bonds formed by hybrid orbitals; delocalized pairs of π-electrons (in the case of multiple bonds) move in their electrostatic field (Table 5.3). sp hybridization. This type of hybridization occurs when an atom forms two bonds due to electrons located in the s- and p-orbitals and having similar energies. This type of hybridization is characteristic of AB2 type molecules (Fig. 5.4). Examples of such molecules and ions are given in table. 5.3 (Fig. 5.4).

Table 5.3

Geometric shapes of molecules

E - lone electron pair.

Structure of the BeCl2 molecule. A beryllium atom in its normal state has two paired s-electrons in its outer layer. As a result of excitation, one of the s electrons goes into the p-state - two unpaired electrons appear, differing in orbital shape and energy. When a chemical bond is formed, they are converted into two identical sp-hybrid orbitals, directed at an angle of 180 degrees to each other.

Be 2s2 Be 2s1 2p1 - excited state of the atom

Rice. 5.4. Spatial arrangement of sp-hybrid clouds

Main types of intermolecular interactions. Substance in a condensed state. Factors that determine the energy of intermolecular interactions. Hydrogen bond. The nature of the hydrogen bond. Quantitative characteristics of hydrogen bonding. Inter- and intramolecular hydrogen bonding.

INTERMOLECULAR INTERACTIONS- interaction molecules between themselves, without leading to rupture or formation of new chemicals. connections. M.v. determines the difference between real gases and ideal gases, the existence of liquids and mol. crystals. From M. v. depend on plural structural, spectral, thermodynamic. and other saints. The emergence of the concept of M. v. is associated with the name of Van der Waals, who in 1873 proposed a level of state that takes into account the magnesium of matter to explain the properties of real gases and liquids. Therefore, the forces of M. v. often called van der Waals.

The basis of M. century. constitute Coulomb forces interaction. between the electrons and nuclei of one molecule and the nuclei and electrons of another. In the experimentally determined properties of the substance, an averaged interaction is manifested, which depends on the distance R between the molecules, their mutual orientation, structure and physical properties. characteristics (dipole moment, polarizability, etc.). At large R, which significantly exceeds the linear dimensions of the molecules themselves, as a result of which the electron shells of the molecules do not overlap, the forces of M.V. can be quite reasonably divided into three types - electrostatic, polarization (induction) and dispersive. Electrostatic forces are sometimes called orientational, but this is inaccurate, since the mutual orientation of molecules can also be determined by polarization. forces if the molecules are anisotropic.

At small distances between molecules (R ~ l), distinguish between individual types of molecules. can only be approximated, and in addition to the three named types, two more are distinguished, related to the overlap of electronic shells - exchange interaction and interactions due to electron charge transfer. Despite a certain convention, such a division in each specific case makes it possible to explain the nature of M. century. and calculate its energy.

The structure of matter in the condensed state.

Depending on the distance between the particles that make up a substance, and on the nature and energy of interaction between them, a substance can be in one of three states of aggregation: solid, liquid and gaseous.

At a sufficiently low temperature, the substance is in a solid state. The distances between the particles of a crystalline substance are of the order of the size of the particles themselves. The average potential energy of particles is greater than their average kinetic energy. The movement of the particles that make up the crystals is very limited. The forces acting between particles keep them in close equilibrium positions. This explains the presence of crystalline bodies with their own shape and volume and high shear resistance.

When melting, solids turn into liquids. In structure, a liquid substance differs from a crystalline one in that not all particles are located at the same distances from each other as in crystals; some molecules are distant from each other at large distances. The average kinetic energy of particles for substances in the liquid state is approximately equal to their average potential energy.

Solid and liquid states are often combined under the common term condensed state.

Types of intermolecular interactions intramolecular hydrogen bond. Bonds in the formation of which the restructuring of electronic shells does not occur are called interaction between molecules . The main types of molecular interactions include van der Waals forces, hydrogen bonds and donor-acceptor interactions.

When molecules come together, attraction appears, which causes the appearance of a condensed state of matter (liquid, solid with a molecular crystal lattice). The forces that promote the attraction of molecules are called van der Waals forces.

They are characterized by three types intermolecular interaction :

a) orientational interaction, which manifests itself between polar molecules tending to occupy a position in which their dipoles would face each other with opposite poles, and the moment vectors of these dipoles would be oriented along the same straight line (in another way it is called dipole-dipole interaction );

b) induction, which arises between induced dipoles, the reason for the formation of which is the mutual polarization of the atoms of two approaching molecules;

c) dispersive, which arises as a result of the interaction of microdipoles formed due to instantaneous displacements of positive and negative charges in molecules during the movement of electrons and vibrations of nuclei.

Dispersion forces act between any particles. Orientational and inductive interactions do not occur for particles of many substances, for example: He, Ar, H2, N2, CH4. For NH3 molecules, dispersion interaction accounts for 50%, orientation interaction accounts for 44.6%, and induction interaction accounts for 5.4%. The polar energy of van der Waals attractive forces is characterized by low values. So, for ice it is 11 kJ/mol, i.e. 2.4% H-O covalent bond energy (456 kJ/mol). Vander Waals forces of attraction are physical interactions.

Hydrogen bond is a physicochemical bond between the hydrogen of one molecule and the EO element of another molecule. The formation of hydrogen bonds is explained by the fact that in polar molecules or groups the polarized hydrogen atom has unique properties: the absence of internal electron shells, a significant shift of the electron pair to an atom with high EO and a very small size. Therefore, hydrogen is able to penetrate deeply into the electron shell of a neighboring negatively polarized atom. As spectral data show, the donor-acceptor interaction of the EO atom as a donor and the hydrogen atom as an acceptor also plays a significant role in the formation of a hydrogen bond. Hydrogen bonding can be intermolecular or intramolecular.

Hydrogen bonds can occur both between different molecules and within a molecule if this molecule contains groups with donor and acceptor abilities. Thus, it is intramolecular hydrogen bonds that play the main role in the formation of peptide chains, which determine the structure of proteins. One of the most famous examples of the influence of intramolecular hydrogen bonding on structure is deoxyribonucleic acid (DNA). The DNA molecule is folded into a double helix. The two strands of this double helix are linked to each other by hydrogen bonds. The hydrogen bond is intermediate in nature between valence and intermolecular interactions. It is associated with the unique properties of the polarized hydrogen atom, its small size and the absence of electronic layers.

Intermolecular and intramolecular hydrogen bonding.

Hydrogen bonds are found in many chemical compounds. They arise, as a rule, between atoms of fluorine, nitrogen and oxygen (the most electronegative elements), less often - with the participation of atoms of chlorine, sulfur and other non-metals. Strong hydrogen bonds are formed in liquid substances such as water, hydrogen fluoride, oxygen-containing inorganic acids, carboxylic acids, phenols, alcohols, ammonia, and amines. During crystallization, hydrogen bonds in these substances are usually preserved. Therefore, their crystal structures take the form of chains (methanol), flat two-dimensional layers (boric acid), or spatial three-dimensional networks (ice).

If a hydrogen bond unites parts of one molecule, then we speak of intramolecular hydrogen bond. This is especially true for many organic compounds (Fig. 42). If a hydrogen bond is formed between a hydrogen atom of one molecule and a non-metal atom of another molecule (intermolecular hydrogen bond), then the molecules form fairly strong pairs, chains, rings. Thus, formic acid exists in the form of dimers in both liquid and gaseous states:

and hydrogen fluoride gas contains polymer molecules containing up to four HF particles. Strong bonds between molecules can be found in water, liquid ammonia, and alcohols. All carbohydrates, proteins, and nucleic acids contain the oxygen and nitrogen atoms necessary for the formation of hydrogen bonds. It is known, for example, that glucose, fructose and sucrose are highly soluble in water. An important role in this is played by hydrogen bonds formed in solution between water molecules and numerous OH groups of carbohydrates.

Periodic law. Modern formulation of the periodic law. The periodic table of chemical elements is a graphic illustration of the periodic law. Modern version of the Periodic Table. Features of filling atomic orbitals with electrons and the formation of periods. s-, p-, d-, f- Elements and their arrangement in the periodic table. Groups, periods. Main and secondary subgroups. Boundaries of the periodic system.

Discovery of the Periodic Law.

The basic law of chemistry - the Periodic Law was discovered by D.I. Mendeleev in 1869 at a time when the atom was considered indivisible and nothing was known about its internal structure. The basis of the Periodic Law D.I. Mendeleev laid down atomic masses (formerly atomic weights) and chemical properties of elements.

Having arranged the 63 elements known at that time in order of increasing atomic masses, D.I. Mendeleev obtained a natural (natural) series of chemical elements, in which he discovered the periodic repeatability of chemical properties.

For example, the properties of the typical metal lithium Li were repeated in the elements sodium Na and potassium K, the properties of the typical nonmetal fluorine F were repeated in the elements chlorine Cl, bromine Br, iodine I.

Some elements have D.I. Mendeleev did not discover chemical analogs (for example, aluminum Al and silicon Si), since such analogs were still unknown at that time. For them, he left empty spaces in the natural series and, based on periodic repetition, predicted their chemical properties. After the discovery of the corresponding elements (an analogue of aluminum - gallium Ga, an analogue of silicon - germanium Ge, etc.), the predictions of D.I. Mendeleev were completely confirmed.

Topics of the Unified State Examination codifier: Covalent chemical bond, its varieties and mechanisms of formation. Characteristics of covalent bonds (polarity and bond energy). Ionic bond. Metal connection. Hydrogen bond

Intramolecular chemical bonds

First, let's look at the bonds that arise between particles within molecules. Such connections are called intramolecular.

Chemical bond between atoms of chemical elements has an electrostatic nature and is formed due to interaction of external (valence) electrons, in more or less degree held by positively charged nuclei bonded atoms.

The key concept here is ELECTRONEGATIVITY. It is this that determines the type of chemical bond between atoms and the properties of this bond.

is the ability of an atom to attract (hold) external(valence) electrons. Electronegativity is determined by the degree of attraction of outer electrons to the nucleus and depends primarily on the radius of the atom and the charge of the nucleus.

Electronegativity is difficult to determine unambiguously. L. Pauling compiled a table of relative electronegativities (based on the bond energies of diatomic molecules). The most electronegative element is fluorine with meaning 4 .

It is important to note that in different sources you can find different scales and tables of electronegativity values. This should not be alarmed, since the formation of a chemical bond plays a role atoms, and it is approximately the same in any system.

If one of the atoms in the A:B chemical bond attracts electrons more strongly, then the electron pair moves towards it. The more electronegativity difference atoms, the more the electron pair shifts.

If the electronegativities of interacting atoms are equal or approximately equal: EO(A)≈EO(B), then the common electron pair does not shift to any of the atoms: A: B. This connection is called covalent nonpolar.

If the electronegativities of the interacting atoms differ, but not greatly (the difference in electronegativity is approximately from 0.4 to 2: 0,4<ΔЭО<2 ), then the electron pair is displaced to one of the atoms. This connection is called covalent polar .

If the electronegativities of interacting atoms differ significantly (the difference in electronegativity is greater than 2: ΔEO>2), then one of the electrons is almost completely transferred to another atom, with the formation ions. This connection is called ionic.

Basic types of chemical bonds − covalent, ionic And metal communications. Let's take a closer look at them.

Covalent chemical bond

Covalent bond – it's a chemical bond , formed due to formation of a common electron pair A:B . Moreover, two atoms overlap atomic orbitals. A covalent bond is formed by the interaction of atoms with a small difference in electronegativity (usually between two non-metals) or atoms of one element.

Basic properties of covalent bonds

• focus,
• saturability,
• polarity,
• polarizability.

These bonding properties influence the chemical and physical properties of substances.

Communication direction characterizes the chemical structure and form of substances. The angles between two bonds are called bond angles. For example, in a water molecule the bond angle H-O-H is 104.45 o, therefore the water molecule is polar, and in a methane molecule the bond angle H-C-H is 108 o 28′.

Saturability is the ability of atoms to form a limited number of covalent chemical bonds. The number of bonds that an atom can form is called.

Polarity bonding occurs due to the uneven distribution of electron density between two atoms with different electronegativity. Covalent bonds are divided into polar and nonpolar.

Polarizability connections are the ability of bond electrons to shift under the influence of an external electric field(in particular, the electric field of another particle). Polarizability depends on electron mobility. The further the electron is from the nucleus, the more mobile it is, and accordingly the molecule is more polarizable.

Covalent nonpolar chemical bond

There are 2 types of covalent bonding – POLAR And NON-POLAR .

Example . Let's consider the structure of the hydrogen molecule H2. Each hydrogen atom in its outer energy level carries 1 unpaired electron. To display an atom, we use the Lewis structure - this is a diagram of the structure of the outer energy level of an atom, when electrons are indicated by dots. Lewis point structure models are quite helpful when working with elements of the second period.

H. + . H = H:H

Thus, a hydrogen molecule has one shared electron pair and one H–H chemical bond. This electron pair does not shift to any of the hydrogen atoms, because Hydrogen atoms have the same electronegativity. This connection is called covalent nonpolar .

Covalent nonpolar (symmetric) bond is a covalent bond formed by atoms with equal electronegativity (usually the same nonmetals) and, therefore, with a uniform distribution of electron density between the nuclei of atoms.

The dipole moment of non-polar bonds is 0.

Examples: H 2 (H-H), O 2 (O=O), S 8.

Covalent polar chemical bond

Covalent polar bond is a covalent bond that occurs between atoms with different electronegativity (usually, various non-metals) and is characterized displacement shared electron pair to a more electronegative atom (polarization).

The electron density is shifted to the more electronegative atom - therefore, a partial negative charge (δ-) appears on it, and a partial positive charge (δ+, delta +) appears on the less electronegative atom.

The greater the difference in electronegativity of atoms, the higher polarity connections and more dipole moment . Additional attractive forces act between neighboring molecules and charges of opposite sign, which increases strength communications.

Bond polarity affects the physical and chemical properties of compounds. The reaction mechanisms and even the reactivity of neighboring bonds depend on the polarity of the bond. The polarity of the connection often determines molecule polarity and thus directly affects such physical properties as boiling point and melting point, solubility in polar solvents.

Examples: HCl, CO 2, NH 3.

Mechanisms of covalent bond formation

Covalent chemical bonds can occur by 2 mechanisms:

1. Exchange mechanism the formation of a covalent chemical bond is when each particle provides one unpaired electron to form a common electron pair:

A . + . B= A:B

2. Covalent bond formation is a mechanism in which one of the particles provides a lone pair of electrons, and the other particle provides a vacant orbital for this electron pair:

A: + B= A:B

In this case, one of the atoms provides a lone pair of electrons ( donor), and the other atom provides a vacant orbital for that pair ( acceptor). As a result of the formation of both bonds, the energy of the electrons decreases, i.e. this is beneficial for the atoms.

A covalent bond formed by a donor-acceptor mechanism is not different in properties from other covalent bonds formed by the exchange mechanism. The formation of a covalent bond by the donor-acceptor mechanism is typical for atoms either with a large number of electrons at the external energy level (electron donors), or, conversely, with a very small number of electrons (electron acceptors). The valence capabilities of atoms are discussed in more detail in the corresponding section.

A covalent bond is formed by a donor-acceptor mechanism:

- in a molecule carbon monoxide CO(the bond in the molecule is triple, 2 bonds are formed by the exchange mechanism, one by the donor-acceptor mechanism): C≡O;

- V ammonium ion NH 4 +, in ions organic amines, for example, in the methylammonium ion CH 3 -NH 2 + ;

- V complex compounds, a chemical bond between the central atom and ligand groups, for example, in sodium tetrahydroxoaluminate Na bond between aluminum and hydroxide ions;

- V nitric acid and its salts- nitrates: HNO 3, NaNO 3, in some other nitrogen compounds;

- in a molecule ozone O3.

Basic characteristics of covalent bonds

Covalent bonds typically form between nonmetal atoms. The main characteristics of a covalent bond are length, energy, multiplicity and directionality.

Multiplicity of chemical bond

Multiplicity of chemical bond - This number of shared electron pairs between two atoms in a compound. The multiplicity of a bond can be determined quite easily from the values ​​of the atoms that form the molecule.

For example , in the hydrogen molecule H 2 the bond multiplicity is 1, because Each hydrogen has only 1 unpaired electron in its outer energy level, hence one shared electron pair is formed.

In the O 2 oxygen molecule, the bond multiplicity is 2, because Each atom at the outer energy level has 2 unpaired electrons: O=O.

In the nitrogen molecule N2, the bond multiplicity is 3, because between each atom there are 3 unpaired electrons at the outer energy level, and the atoms form 3 common electron pairs N≡N.

Covalent bond length

Chemical bond length is the distance between the centers of the nuclei of the atoms forming the bond. It is determined by experimental physical methods. The bond length can be estimated approximately using the additivity rule, according to which the bond length in the AB molecule is approximately equal to half the sum of the bond lengths in molecules A 2 and B 2:

The length of a chemical bond can be roughly estimated by atomic radii forming a bond, or by communication multiplicity, if the radii of the atoms are not very different.

As the radii of the atoms forming a bond increase, the bond length will increase.

For example

As the multiplicity of bonds between atoms increases (the atomic radii of which do not differ or differ only slightly), the bond length will decrease.

For example . In the series: C–C, C=C, C≡C, the bond length decreases.

Communication energy

A measure of the strength of a chemical bond is the bond energy. Communication energy determined by the energy required to break a bond and remove the atoms forming that bond to an infinitely large distance from each other.

A covalent bond is very durable. Its energy ranges from several tens to several hundred kJ/mol. The higher the bond energy, the greater the bond strength, and vice versa.

The strength of a chemical bond depends on the bond length, bond polarity, and bond multiplicity. The longer a chemical bond, the easier it is to break, and the lower the bond energy, the lower its strength. The shorter the chemical bond, the stronger it is, and the greater the bond energy.

For example, in the series of compounds HF, HCl, HBr from left to right, the strength of the chemical bond decreases, because The connection length increases.

Ionic chemical bond

Ionic bond is a chemical bond based on electrostatic attraction of ions.

Ions are formed in the process of accepting or donating electrons by atoms. For example, atoms of all metals weakly hold electrons from the outer energy level. Therefore, metal atoms are characterized by restorative properties- ability to donate electrons.

Example. The sodium atom contains 1 electron at energy level 3. By easily giving it up, the sodium atom forms the much more stable Na + ion, with the electron configuration of the noble gas neon Ne. The sodium ion contains 11 protons and only 10 electrons, so the total charge of the ion is -10+11 = +1:

+11Na) 2 ) 8 ) 1 - 1e = +11 Na +) 2 ) 8

Example. A chlorine atom in its outer energy level contains 7 electrons. To acquire the configuration of a stable inert argon atom Ar, chlorine needs to gain 1 electron. After adding an electron, a stable chlorine ion is formed, consisting of electrons. The total charge of the ion is -1:

+17Cl) 2 ) 8 ) 7 + 1e = +17 Cl — ) 2 ) 8 ) 8

Note:

• The properties of ions are different from the properties of atoms!
• Stable ions can form not only atoms, but also groups of atoms. For example: ammonium ion NH 4 +, sulfate ion SO 4 2-, etc. Chemical bonds formed by such ions are also considered ionic;
• Ionic bonds are usually formed between each other metals And nonmetals(non-metal groups);

The resulting ions are attracted due to electrical attraction: Na + Cl -, Na 2 + SO 4 2-.

Let us visually summarize difference between covalent and ionic bond types:

Metal connection is a connection that is formed relatively free electrons between metal ions, forming a crystal lattice.

Metal atoms are usually located on the outer energy level one to three electrons. The radii of metal atoms, as a rule, are large - therefore, metal atoms, unlike non-metals, give up their outer electrons quite easily, i.e. are strong reducing agents.

By donating electrons, metal atoms turn into positively charged ions . The detached electrons are relatively free are moving between positively charged metal ions. Between these particles a connection arises, because shared electrons hold metal cations arranged in layers together , thus creating a fairly strong metal crystal lattice . In this case, the electrons continuously move chaotically, i.e. New neutral atoms and new cations constantly appear.

Intermolecular interactions

Separately, it is worth considering the interactions that arise between individual molecules in a substance - intermolecular interactions . Intermolecular interactions are a type of interaction between neutral atoms in which no new covalent bonds appear. The forces of interaction between molecules were discovered by Van der Waals in 1869, and named after him Van dar Waals forces. Van der Waals forces are divided into orientation, induction And dispersive . The energy of intermolecular interactions is much less than the energy of chemical bonds.

Orientation forces of attraction occur between polar molecules (dipole-dipole interaction). These forces occur between polar molecules. Inductive interactions is the interaction between a polar molecule and a non-polar one. A nonpolar molecule is polarized due to the action of a polar one, which generates additional electrostatic attraction.

A special type of intermolecular interaction is hydrogen bonds. - these are intermolecular (or intramolecular) chemical bonds that arise between molecules that have highly polar covalent bonds - H-F, H-O or H-N. If there are such bonds in a molecule, then between the molecules there will be additional attractive forces .

Education mechanism hydrogen bonding is partly electrostatic and partly donor-acceptor. In this case, the electron pair donor is an atom of a strongly electronegative element (F, O, N), and the acceptor is the hydrogen atoms connected to these atoms. Hydrogen bonds are characterized by focus in space and saturation

Hydrogen bonds can be indicated by dots: H ··· O. The greater the electronegativity of the atom connected to hydrogen, and the smaller its size, the stronger the hydrogen bond. It is typical primarily for connections fluorine with hydrogen , as well as to oxygen and hydrogen , less nitrogen with hydrogen .

Hydrogen bonds occur between the following substances:

— hydrogen fluoride HF(gas, solution of hydrogen fluoride in water - hydrofluoric acid), water H 2 O (steam, ice, liquid water):

— solution of ammonia and organic amines- between ammonia and water molecules;

— organic compounds in which O-H or N-H bonds: alcohols, carboxylic acids, amines, amino acids, phenols, aniline and its derivatives, proteins, solutions of carbohydrates - monosaccharides and disaccharides.

Hydrogen bonding affects the physical and chemical properties of substances. Thus, additional attraction between molecules makes it difficult for substances to boil. Substances with hydrogen bonds exhibit an abnormal increase in boiling point.

For example As a rule, with increasing molecular weight, an increase in the boiling point of substances is observed. However, in a number of substances H 2 O-H 2 S-H 2 Se-H 2 Te we do not observe a linear change in boiling points.

Namely, at water boiling point is abnormally high - no less than -61 o C, as the straight line shows us, but much more, +100 o C. This anomaly is explained by the presence of hydrogen bonds between water molecules. Therefore, under normal conditions (0-20 o C) water is liquid by phase state.