Topics of the USE codifier: Covalent chemical bond, its varieties and mechanisms of formation. Characteristics of a covalent bond (polarity and bond energy). Ionic bond. Metal connection. hydrogen bond

Intramolecular chemical bonds

Let us first consider 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 interactions of external (valence) electrons, in more or less degree held by positively charged nuclei bonded atoms.

The key concept here is ELECTRONEGNATIVITY. It is she who 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 external electrons to the nucleus and depends mainly 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 electronegativity (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 frightened, 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 chemical bond A:B attracts electrons more strongly, then the electron pair is shifted towards it. The more electronegativity difference atoms, the more the electron pair is displaced.

If the electronegativity values ​​of the interacting atoms are equal or approximately equal: EO(A)≈EO(V), then the shared electron pair is not displaced to any of the atoms: A: B. Such a connection is called covalent non-polar.

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

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

The main types of chemical bonds are − covalent, ionic and metallic connections. Let's consider them in more detail.

covalent chemical bond

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

Basic properties of covalent bonds

• orientation,
• saturability,
• polarity,
• polarizability.

These bond properties affect the chemical and physical properties of substances.

Direction of communication characterizes the chemical structure and form of substances. The angles between two bonds are called bond angles. For example, in a water molecule, the H-O-H bond angle is 104.45 o, so the water molecule is polar, and in the methane molecule, the H-C-H bond angle 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 bonds arise due to the uneven distribution of electron density between two atoms with different electronegativity. Covalent bonds are divided into polar and non-polar.

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

Covalent non-polar chemical bond

There are 2 types of covalent bonding - POLAR and NON-POLAR .

Example . Consider the structure of the hydrogen molecule H 2 . Each hydrogen atom carries 1 unpaired electron in its outer energy level. To display an atom, we use the Lewis structure - this is a diagram of the structure of the external energy level of an atom, when electrons are denoted by dots. Lewis point structure models are a good help when working with elements of the second period.

H. + . H=H:H

Thus, the hydrogen molecule has one common electron pair and one H–H chemical bond. This electron pair is not displaced to any of the hydrogen atoms, because the electronegativity of hydrogen atoms is the same. Such a connection is called covalent non-polar .

Covalent non-polar (symmetrical) bond - this is a covalent bond formed by atoms with equal electronegativity (as a rule, the same non-metals) and, therefore, with a uniform distribution of electron density between the nuclei of atoms.

The dipole moment of nonpolar 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, different non-metals) and is characterized displacement common electron pair to a more electronegative atom (polarization).

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

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

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 a bond often determines polarity of the molecule and thus directly affects such physical properties as boiling point and melting point, solubility in polar solvents.

Examples: HCl, CO 2 , NH 3 .

Mechanisms for the formation of a covalent bond

A covalent chemical bond can occur by 2 mechanisms:

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

BUT . + . B= A:B

2. The formation of a covalent bond is such a mechanism in which one of the particles provides an unshared electron pair, and the other particle provides a vacant orbital for this electron pair:

BUT: + B= A:B

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

A covalent bond formed by the donor-acceptor mechanism, is not different by 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 in the external energy level (electron donors), or vice versa, with a very small number of electrons (electron acceptors). The valence possibilities of atoms are considered in more detail in the corresponding.

A covalent bond is formed by the 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;

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

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

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

- in a molecule ozone O 3 .

Main characteristics of a covalent bond

A covalent bond, as a rule, is formed between the atoms of non-metals. The main characteristics of a covalent bond are length, energy, multiplicity and directivity.

Chemical bond multiplicity

Chemical bond multiplicity - This the number of shared electron pairs between two atoms in a compound. The multiplicity of the bond can be quite easily determined from the value 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 the outer energy level, therefore, one common electron pair is formed.

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

In the nitrogen molecule N 2, the bond multiplicity is 3, because between each atom there are 3 unpaired electrons in 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 atoms that form a bond. It is determined by experimental physical methods. The bond length can be estimated approximately, according to the additivity rule, according to which the bond length in the AB molecule is approximately equal to half the sum of the bond lengths in the A 2 and B 2 molecules:

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

With an increase in the radii of the atoms forming a bond, the bond length will increase.

for example

With an increase in the multiplicity of bonds between atoms (whose atomic radii do not differ, or differ slightly), the bond length will decrease.

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

Bond energy

A measure of the strength of a chemical bond is the bond energy. Bond energy is determined by the energy required to break the bond and remove the atoms that form this bond to an infinite distance from each other.

The covalent bond is very durable. Its energy ranges from several tens to several hundreds of kJ/mol. The greater 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 the 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 length of the bond 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 giving away electrons by atoms. For example, the atoms of all metals weakly hold the electrons of the outer energy level. Therefore, metal atoms are characterized restorative properties the ability to donate electrons.

Example. The sodium atom contains 1 electron at the 3rd energy level. Easily giving it away, the sodium atom forms a much more stable Na + ion, with the electron configuration of the noble neon gas 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. The chlorine atom has 7 electrons in its outer energy level. To acquire the configuration of a stable inert argon atom Ar, chlorine needs to attach 1 electron. After the attachment of 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 metals and nonmetals(groups of non-metals);

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

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

metal connection is the relationship that is formed relatively free electrons between metal ions forming a crystal lattice.

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

By donating electrons, metal atoms become positively charged ions . The detached electrons are relatively free are moving between positively charged metal ions. Between these particles there is a connection, because shared electrons hold metal cations in layers together , thus creating a sufficiently strong metal crystal lattice . In this case, the electrons continuously move randomly, i.e. new neutral atoms and new cations are constantly emerging.

Intermolecular interactions

Separately, it is worth considering the interactions that occur between individual molecules in a substance - intermolecular interactions . Intermolecular interactions are a type of interaction between neutral atoms in which new covalent bonds do not 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 dispersion . The energy of intermolecular interactions is much less than the energy of a chemical bond.

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

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

Mechanism of Education The hydrogen bond is partly electrostatic and partly donor-acceptor. In this case, an atom of a strongly electronegative element (F, O, N) acts as an electron pair donor, and hydrogen atoms connected to these atoms act as an acceptor. Hydrogen bonds are characterized orientation in space and saturation .

The hydrogen bond can be denoted by dots: H ··· O. The greater the electronegativity of an atom connected to hydrogen, and the smaller its size, the stronger the hydrogen bond. It is primarily characteristic of compounds fluorine with hydrogen , as well as to oxygen with 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.

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

for example As a rule, with an increase in 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 boiling point of water is abnormally high - not 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.

There is no unified theory of chemical bonding; chemical bonding is conditionally divided into covalent (universal type of bond), ionic (a special case of 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 exchange mechanism the formation of a covalent bond occurs due to the socialization of common electron pairs. In this case, each atom tends to acquire an inert gas shell, i.e. get the completed outer energy level. The formation of an exchange-type chemical bond 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 the structure of the atom and quantum mechanics, the formation of a covalent bond is represented as an 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 its energy. A covalent bond can be formed by overlapping different orbitals. As a result of the overlapping of s-s, s-p orbitals, as well as d-d, p-p, d-p orbitals by the side lobes, a bond is formed. Perpendicular to the line connecting the nuclei of 2 atoms, a bond is formed. One - and one - bonds are able to form a multiple (double) covalent bond, characteristic of organic substances of the 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).

The formation of a covalent bond donor-acceptor mechanism consider 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 the formation of a covalent bond 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 covalent bond characteristics 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); directivity (this value is associated with the geometry of the molecule and the concept of "valence angle" - the angle between bonds).

Ionic bond

There are no compounds with a pure ionic bond, although this is understood as such a chemically bound state of atoms in which a stable electronic environment of the atom is created with the complete transition of the total electron density to an 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 called electrically charged particles formed by detaching or attaching an electron to an atom.

When transferring an electron, the atoms of metals and non-metals tend to form a stable configuration of the electron shell 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 an ionic bond exists in its pure form are found in the vapor state of a substance. The ionic bond is very strong, in connection with this, substances with this bond have a high melting point. Unlike covalent bonds, ionic bonds are not characterized by directivity and saturation, since the electric field created by ions acts equally on all ions due to spherical symmetry.

metal bond

A metallic bond is realized only in metals - this is an interaction that holds metal atoms in a single lattice. Only the valence electrons of the metal atoms, which belong to its entire volume, participate in the formation of the bond. In metals, electrons are constantly detached from atoms, which move throughout the mass of the metal. Metal atoms, devoid of electrons, turn into positively charged ions, which tend to take moving electrons towards them. 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, therefore, metals are characterized by a high melting point, and the presence of an "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 introduced 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 by the donor-acceptor mechanism.

The concept of a chemical bond is of no small importance in various fields of chemistry as a science. This is due to the fact that it is with its help that individual atoms are able to combine into molecules, forming all kinds of substances, which, in turn, are the subject of chemical research.

The variety of atoms and molecules is associated with the emergence of various types of bonds between them. Different classes of molecules are characterized by their own features of the distribution of electrons, and hence their own types of bonds.

Basic concepts

chemical bond called a set of interactions that lead to the binding of atoms to form stable particles of a more complex structure (molecules, ions, radicals), as well as aggregates (crystals, glasses, etc.). The nature of these interactions is electrical in nature, and they arise during the distribution of valence electrons in approaching atoms.

Valency accepted name the ability of an atom to form a certain number of bonds with other atoms. In ionic compounds, the number of given or attached electrons is taken as the value of valency. In covalent compounds, it is equal to the number of common electron pairs.

Under the degree of oxidation is understood as conditional the charge that could be on an atom if all polar covalent bonds were ionic.

The multiplicity of the connection is called the number of shared electron pairs between the considered atoms.

The bonds considered in various branches of chemistry can be divided into two types of chemical bonds: those that lead to the formation of new substances (intramolecular) , and those that arise between molecules (intermolecular).

Basic communication characteristics

By bond energy is the energy required to break all the bonds in a molecule. It is also the energy released during bond formation.

Communication length called such a distance between neighboring nuclei of atoms in a molecule, at which the forces of attraction and repulsion are balanced.

These two characteristics of the chemical bond of atoms are a measure of its strength: the shorter the length and the greater the energy, the stronger the bond.

Valence angle It is customary to call the angle between the represented lines passing in the direction of the bond through the nuclei of atoms.

Relationship Description Methods

The two most common approaches to explaining the chemical bond, borrowed from quantum mechanics:

Method of molecular orbitals. He considers a molecule as a set of electrons and nuclei of atoms, with each individual electron moving in the field of action of all other electrons and nuclei. The molecule has an orbital structure, and all its electrons are distributed along these orbits. Also, this method is called MO LCAO, which stands for "molecular orbital - linear combination

The method of valence bonds. Represents a molecule as a system of two central molecular orbitals. Moreover, each of them corresponds to one bond between two adjacent atoms in the molecule. The method is based on the following provisions:

1. The formation of a chemical bond is carried out by a pair of electrons with opposite spins, which are located between the two considered atoms. The formed electron pair belongs to two atoms equally.
2. The number of bonds formed by one or another atom is equal to the number of unpaired electrons in the ground and excited states.
3. If electron pairs do not take part in the formation of a bond, then they are called lone pairs.

Electronegativity

The type of chemical bond in substances can be determined based on the difference in the electronegativity values ​​of its constituent atoms. Under electronegativity understand the ability of atoms to attract common electron pairs (electron cloud), which leads to bond polarization.

There are various ways to determine the values ​​of the electronegativity of chemical elements. However, the most used is the scale based on thermodynamic data, which was proposed back in 1932 by L. Pauling.

The greater the difference in the electronegativity of atoms, the more pronounced its ionicity. On the contrary, equal or close electronegativity values ​​indicate the covalent nature of the bond. In other words, it is possible to determine which chemical bond is observed in a particular molecule mathematically. To do this, you need to calculate ΔX - the difference in the electronegativity of atoms according to the formula: ΔX=|X 1 -X 2 |.

• If a ΔX>1.7, then the bond is ionic.
• If a 0.5≤ΔХ≤1.7, the covalent bond is polar.
• If a ΔX=0 or close to it, then the bond is covalent non-polar.

Ionic bond

An ionic bond is such a bond that appears between ions or due to the complete withdrawal of a common electron pair by one of the atoms. In substances, this type of chemical bonding is carried out by the forces of electrostatic attraction.

Ions are charged particles formed from atoms as a result of the addition or release of electrons. When an atom accepts electrons, it acquires a negative charge and becomes an anion. If an atom donates valence electrons, it becomes a positively charged particle called a cation.

It is characteristic of compounds formed by the interaction of atoms of typical metals with atoms of typical non-metals. The main of this process is the aspiration of atoms to acquire stable electronic configurations. And for this, typical metals and non-metals need to give or accept only 1-2 electrons, which they do with ease.

The mechanism of formation of an ionic chemical bond in a molecule is traditionally considered using the example of the interaction of sodium and chlorine. Alkali metal atoms easily donate an electron pulled by a halogen atom. As a result, the Na + cation and the Cl - anion are formed, which are held together by electrostatic attraction.

There is no ideal ionic bond. Even in such compounds, which are often referred to as ionic, the final transfer of electrons from atom to atom does not occur. The formed electron pair still remains in common use. Therefore, they talk about the degree of ionicity of a covalent bond.

An ionic bond is characterized by two main properties related to each other:

• non-directionality, i.e., the electric field around the ion has the shape of a sphere;
• unsaturation, i.e., the number of oppositely charged ions that can be placed around any ion, is determined by their size.

covalent chemical bond

The bond formed when the electron clouds of non-metal atoms overlap, that is, carried out by a common electron pair, is called a covalent bond. The number of shared pairs of electrons determines the multiplicity of the bond. So, hydrogen atoms are connected by a single H··H bond, and oxygen atoms form a double bond O::O.

There are two mechanisms for its formation:

• Exchange - each atom represents for the formation of a common pair of one electron: A + B = A: B, while the connection involves external atomic orbitals, on which one electron is located.
• Donor-acceptor - to form a bond, one of the atoms (donor) provides a pair of electrons, and the second (acceptor) provides a free orbital for its placement: A +: B \u003d A: B.

The ways in which electron clouds overlap during the formation of a covalent chemical bond are also different.

1. Direct. The cloud overlap region lies on a straight imaginary line connecting the nuclei of the considered atoms. In this case, σ-bonds are formed. The type of chemical bond that occurs in this case depends on the type of electron clouds undergoing overlap: s-s, s-p, p-p, s-d or p-d σ-bonds. In a particle (molecule or ion), only one σ-bond can occur between two neighboring atoms.
2. Lateral. It is carried out on both sides of the line connecting the nuclei of atoms. This is how a π-bond is formed, and its varieties are also possible: p-p, p-d, d-d. Apart from the σ-bond, the π-bond is never formed; it can be in molecules containing multiple (double and triple) bonds.

Properties of a covalent bond

It is they who determine the chemical and physical characteristics of compounds. The main properties of any chemical bond in substances are its directionality, polarity and polarizability, as well as saturation.

Orientation connections are due to the features of the molecular structure of substances and the geometric shape of their molecules. Its essence lies in the fact that the best overlap of electron clouds is possible with a certain orientation in space. The options for the formation of σ- and π-bonds have already been considered above.

Under satiety understand the ability of atoms to form a certain number of chemical bonds in a molecule. The number of covalent bonds for each atom is limited by the number of outer orbitals.

Polarity bond depends on the difference in the electronegativity values ​​of the atoms. It determines the uniformity of the distribution of electrons between the nuclei of atoms. The covalent bond on this basis can be polar or non-polar.

• If a common electron pair equally belongs to each of the atoms and is located at the same distance from their nuclei, then the covalent bond is non-polar.
• If the common pair of electrons is displaced to the nucleus of one of the atoms, then a covalent polar chemical bond is formed.

Polarizability is expressed by the displacement of bond electrons under the action of an external electric field, which may belong to another particle, neighboring bonds in the same molecule, or come from external sources of electromagnetic fields. Thus, a covalent bond under their influence can change its polarity.

Hybridization of orbitals is understood as a change in their forms during the implementation of a chemical bond. This is necessary to achieve the most effective overlap. There are the following types of hybridization:

• sp3. One s- and three p-orbitals form four "hybrid" orbitals of the same shape. Outwardly, it resembles a tetrahedron with an angle between the axes of 109 °.
• sp2. One s- and two p-orbitals form a flat triangle with an angle between the axes of 120°.
• sp. One s- and one p-orbital form two "hybrid" orbitals with an angle between their axes of 180°.

A feature of the structure of metal atoms is a rather large radius and the presence of a small number of electrons in outer orbitals. As a result, in such chemical elements, the bond between the nucleus and valence electrons is relatively weak and easily broken.

metal a bond is such an interaction between metal atoms-ions, which is carried out with the help of delocalized electrons.

In metal particles, valence electrons can easily leave outer orbitals, as well as occupy vacant places on them. Thus, at different times, the same particle can be an atom and an ion. The electrons torn off from them move freely throughout the entire volume of the crystal lattice and carry out a chemical bond.

This type of bond has similarities with ionic and covalent bonds. As well as for ionic, ions are necessary for the existence of a metallic bond. But if for the implementation of electrostatic interaction in the first case, cations and anions are needed, then in the second, the role of negatively charged particles is played by electrons. If we compare a metallic bond with a covalent bond, then the formation of both requires common electrons. However, unlike a polar chemical bond, they are not localized between two atoms, but belong to all metal particles in the crystal lattice.

The metallic bond is responsible for the special properties of almost all metals:

• plasticity, present due to the possibility of displacement of layers of atoms in the crystal lattice held by the electron gas;
• metallic luster, which is observed due to the reflection of light rays from electrons (in the powder state there is no crystal lattice and, therefore, electrons moving along it);
• electrical conductivity, which is carried out by a stream of charged particles, and in this case, small electrons move freely among large metal ions;
• thermal conductivity is observed due to the ability of electrons to transfer heat.

This type of chemical bond is sometimes referred to as intermediate between covalent and intermolecular interactions. If a hydrogen atom has a bond with one of the strongly electronegative elements (such as phosphorus, oxygen, chlorine, nitrogen), then it is able to form an additional bond, called hydrogen.

It is much weaker than all the types of bonds considered above (the energy is not more than 40 kJ/mol), but it cannot be neglected. That is why the hydrogen chemical bond in the diagram looks like a dotted line.

The occurrence of a hydrogen bond is possible due to the donor-acceptor electrostatic interaction simultaneously. A large difference in the values ​​of electronegativity leads to the appearance of excess electron density on the atoms O, N, F and others, as well as its lack on the hydrogen atom. In the event that there is no existing chemical bond between such atoms, attractive forces are activated if they are close enough. In this case, the proton is an electron pair acceptor, and the second atom is a donor.

A hydrogen bond can occur both between neighboring molecules, for example, water, carboxylic acids, alcohols, ammonia, and within a molecule, for example, salicylic acid.

The presence of a hydrogen bond between water molecules explains a number of its unique physical properties:

• The values ​​of its heat capacity, dielectric constant, boiling and melting points, in accordance with the calculations, should be much lower than the real ones, which is explained by the bonding of molecules and the need to expend energy to break intermolecular hydrogen bonds.
• Unlike other substances, as the temperature decreases, the volume of water increases. This is due to the fact that the molecules occupy a certain position in the crystal structure of ice and move away from each other by the length of the hydrogen bond.

This bond plays a special role for living organisms, since its presence in protein molecules determines their special structure, and hence their properties. In addition, nucleic acids, making up the double helix of DNA, are also connected precisely by hydrogen bonds.

Bonds in crystals

The vast majority of solids have a crystal lattice - a special mutual arrangement of the particles that form them. In this case, three-dimensional periodicity is observed, and atoms, molecules or ions are located at the nodes, which are connected by imaginary lines. Depending on the nature of these particles and the bonds between them, all crystal structures are divided into atomic, molecular, ionic and metallic.

At the nodes of the ionic crystal lattice are cations and anions. Moreover, each of them is surrounded by a strictly defined number of ions with only the opposite charge. A typical example is sodium chloride (NaCl). They tend to have high melting points and hardness, as they require a lot of energy to break down.

At the nodes of the molecular crystal lattice, there are molecules of substances formed by a covalent bond (for example, I 2). They are connected to each other by a weak van der Waals interaction, and therefore, such a structure is easy to destroy. Such compounds have low boiling and melting points.

The atomic crystal lattice is formed by atoms of chemical elements with high valence values. They are connected by strong covalent bonds, which means that the substances have high boiling and melting points and high hardness. An example is a diamond.

Thus, all types of bonds present in chemical substances have their own characteristics, which explain the intricacies of the interaction of particles in molecules and substances. The properties of the compounds depend on them. They determine all processes occurring in the environment.

Any interaction between atoms is possible only in the presence of a chemical bond. Such a connection is the reason for the formation of a stable polyatomic system - a molecular ion, a molecule, a crystal lattice. A strong chemical bond requires a lot of energy to break, which is why it is the base value for measuring bond strength.

Conditions for the formation of a chemical bond

The formation of a chemical bond is always accompanied by the release of energy. This process occurs due to a decrease in the potential energy of a system of interacting particles - molecules, ions, atoms. The potential energy of the resulting system of interacting elements is always less than the energy of unbound outgoing particles. Thus, the basis for the occurrence of a chemical bond in the system is the decline in the potential energy of its elements.

The nature of the chemical interaction

A chemical bond is a consequence of the interaction of electromagnetic fields that arise around the electrons and nuclei of atoms of those substances that take part in the formation of a new molecule or crystal. After the discovery of the theory of the structure of the atom, the nature of this interaction became more accessible for study.

For the first time, the idea of ​​​​the electrical nature of a chemical bond arose from the English physicist G. Davy, who suggested that molecules are formed due to the electrical attraction of oppositely charged particles. This idea interested the Swedish chemist and naturalist I.Ya. Berzellius, who developed the electrochemical theory of the formation of a chemical bond.

The first theory, which explained the processes of chemical interaction of substances, was imperfect, and over time it had to be abandoned.

Butlerov's theory

A more successful attempt to explain the nature of the chemical bond of substances was made by the Russian scientist A.M. Butlerov. This scientist based his theory on the following assumptions:

• Atoms in the connected state are connected to each other in a certain order. A change in this order causes the formation of a new substance.
• Atoms bind to each other according to the laws of valence.
• The properties of a substance depend on the order of connection of atoms in a molecule of a substance. A different arrangement causes a change in the chemical properties of the substance.
• Atoms bonded together have the strongest influence on each other.

Butlerov's theory explained the properties of chemical substances not only by their composition, but also by the arrangement of atoms. Such an internal order of A.M. Butlerov called "chemical structure".

The theory of the Russian scientist made it possible to put things in order in the classification of substances and made it possible to determine the structure of molecules by their chemical properties. The theory also gave an answer to the question: why molecules containing the same number of atoms have different chemical properties.

Prerequisites for the creation of chemical bond theories

In his theory of the chemical structure, Butlerov did not touch on the question of what a chemical bond is. For this, then there were too few data on the internal structure of matter. Only after the discovery of the planetary model of the atom, the American scientist Lewis began to develop a hypothesis that a chemical bond arises through the formation of an electron pair, which simultaneously belongs to two atoms. Subsequently, this idea became the foundation for the development of the theory of covalent bonds.

covalent chemical bond

A stable chemical compound can be formed when the electron clouds of two neighboring atoms overlap. The result of such mutual crossing is an increasing electron density in the internuclear space. The nuclei of atoms, as you know, are positively charged, and therefore they try to be attracted as close as possible to the negatively charged electron cloud. This attraction is much stronger than the repulsive forces between two positively charged nuclei, so this bond is stable.

The first chemical bond calculations were performed by the chemists Heitler and London. They considered the bond between two hydrogen atoms. The simplest visual representation of it might look like this:

As can be seen, the electron pair occupies a quantum place in both hydrogen atoms. This two-center arrangement of electrons is called a "covalent chemical bond". A covalent bond is typical for molecules of simple substances and their compounds of non-metals. Substances created as a result of a covalent bond usually do not conduct electricity or are semiconductors.

Ionic bond

An ionic type chemical bond occurs when two oppositely charged ions are attracted electrically. Ions can be simple, consisting of one atom of a substance. In compounds of this type, simple ions are most often positively charged atoms of metals of the 1.2 group that have lost their electron. The formation of negative ions is inherent in the atoms of typical non-metals and bases of their acids. Therefore, among the typical ionic compounds, there are many alkali metal halides, such as CsF, NaCl, and others.

Unlike a covalent bond, an ion does not have saturation: a different number of oppositely charged ions can join an ion or group of ions. The number of attached particles is limited only by the linear dimensions of the interacting ions, as well as the condition under which the attractive forces of oppositely charged ions must be greater than the repulsive forces of identically charged particles participating in an ionic type connection.

hydrogen bond

Even before the creation of the theory of chemical structure, it was experimentally observed that hydrogen compounds with various non-metals have somewhat unusual properties. For example, the boiling points of hydrogen fluoride and water are much higher than might be expected.

These and other features of hydrogen compounds can be explained by the ability of the H + atom to form another chemical bond. This type of connection is called a "hydrogen bond". The causes of hydrogen bonding lie in the properties of electrostatic forces. For example, in a hydrogen fluoride molecule, the general electron cloud is so shifted towards fluorine that the space around the atom of this substance is saturated with a negative electric field. Around the hydrogen atom, deprived of its only electron, the field is much weaker and has a positive charge. As a result, there is an additional relationship between the positive fields of electron clouds H + and negative F - .

Chemical bonding of metals

The atoms of all metals are located in space in a certain way. The arrangement of metal atoms is called the crystal lattice. In this case, the electrons of different atoms weakly interact with each other, forming a common electron cloud. This type of interaction between atoms and electrons is called a "metal bond".

It is the free movement of electrons in metals that can explain the physical properties of metallic substances: electrical conductivity, thermal conductivity, strength, fusibility, and others.

3.3.1 Covalent bond - This is a two-center two-electron bond formed due to the overlap of electron clouds carrying unpaired electrons with antiparallel spins. As a rule, it is formed between atoms of one chemical element.

Quantitatively, it is characterized by valency. Element valency - this is its ability to form a certain number of chemical bonds due to free electrons located in the atomic valence zone.

A covalent bond is formed only by a pair of electrons located between atoms. It is called a divided pair. The remaining pairs of electrons are called lone pairs. They fill the shells and do not take part in binding. Communication between atoms can be carried out not only by one, but also by two or even three shared pairs. Such connections are called double and t swarm - multiple bonds.

3.3.1.1 Covalent non-polar bond. A bond carried out by the formation of electron pairs equally belonging to both atoms is called covalent non-polar. It arises between atoms with practically equal electronegativity (0.4 > ΔEO > 0) and, consequently, a uniform distribution of electron density between the nuclei of atoms in homonuclear molecules. For example, H 2 , O 2 , N 2 , Cl 2 , etc. The dipole moment of such bonds is zero. The CH bond in saturated hydrocarbons (for example, in CH 4) is considered practically non-polar, because ΔEO = 2.5 (C) - 2.1 (H) = 0.4.

3.3.1.2 Covalent polar bond. If a molecule is formed by two different atoms, then the overlap zone of electron clouds (orbitals) shifts towards one of the atoms, and such a bond is called polar . With such a connection, the probability of finding electrons near the nucleus of one of the atoms is higher. For example, HCl, H 2 S, PH 3.

Polar (asymmetric) covalent bond - connection between atoms with different electronegativity (2 > ΔEO > 0.4) and asymmetric distribution of a common electron pair. As a rule, it is formed between two non-metals.

The electron density of such a bond is shifted towards a more electronegative atom, which leads to the appearance on it of a partial negative charge  (delta minus), and on a less electronegative atom - a partial positive charge  (delta plus)

C  - Cl

The direction of electron displacement is also indicated by an arrow:

CCl, CO, CN, OH, CMg.

The greater the difference in the electronegativity of the bonded atoms, the higher the polarity of the bond and the greater its dipole moment. Additional forces of attraction act between partial charges of opposite sign. Therefore, the more polar the bond, the stronger it is.

Except polarizability covalent bond has the property satiety - the ability of an atom to form as many covalent bonds as it has energetically available atomic orbitals. The third property of a covalent bond is its orientation.

3.3.2 Ionic bond. The driving force behind its formation is the same aspiration of atoms to the octet shell. But in a number of cases, such an “octet” shell can arise only when electrons are transferred from one atom to another. Therefore, as a rule, an ionic bond is formed between a metal and a non-metal.

Consider as an example the reaction between sodium (3s 1) and fluorine (2s 2 3s 5) atoms. Electronegativity difference in NaF compound

EO = 4.0 - 0.93 = 3.07

Sodium, having donated its 3s 1 electron to fluorine, becomes the Na + ion and remains with a filled 2s 2 2p 6 shell, which corresponds to the electronic configuration of the neon atom. Exactly the same electronic configuration is acquired by fluorine, having accepted one electron donated by sodium. As a result, electrostatic attraction forces arise between oppositely charged ions.

Ionic bond - an extreme case of a polar covalent bond, based on the electrostatic attraction of ions. Such a bond occurs when there is a large difference in the electronegativity of the bonded atoms (EO > 2), when a less electronegative atom almost completely gives up its valence electrons and turns into a cation, and another, more electronegative atom, attaches these electrons and becomes an anion. The interaction of ions of the opposite sign does not depend on the direction, and the Coulomb forces do not have the property of saturation. Because of this ionic bond has no space focus and satiety , since each ion is associated with a certain number of counterions (coordination number of the ion). Therefore, ionically bound compounds do not have a molecular structure and are solid substances that form ionic crystal lattices, with high melting and boiling points, they are highly polar, often salt-like, and electrically conductive in aqueous solutions. For example, MgS, NaCl, A 2 O 3. Compounds with purely ionic bonds practically do not exist, since there is always a certain amount of covalence due to the fact that a complete transition of one electron to another atom is not observed; in the most "ionic" substances, the proportion of bond ionicity does not exceed 90%. For example, in NaF, the bond polarization is about 80%.

In organic compounds, ionic bonds are quite rare, because. a carbon atom tends to neither lose nor gain electrons to form ions.

Valence elements in compounds with ionic bonds very often characterize oxidation state , which, in turn, corresponds to the charge of the ion of the element in the given compound.

Oxidation state is the conditional charge that an atom acquires as a result of the redistribution of electron density. Quantitatively, it is characterized by the number of electrons displaced from a less electronegative element to a more electronegative one. A positively charged ion is formed from the element that gave up its electrons, and a negative ion is formed from the element that received these electrons.

The element in highest oxidation state (maximally positive), has already given up all its valence electrons in the ABD. And since their number is determined by the number of the group in which the element is located, then highest oxidation state for most elements and will be equal to group number . Concerning lowest oxidation state (maximally negative), then it appears during the formation of an eight-electron shell, that is, in the case when the AVZ is completely filled. For non-metals it is calculated according to the formula group number - 8 . For metals is equal to zero because they cannot accept electrons.

For example, the AVZ of sulfur has the form: 3s 2 3p 4 . If an atom gives up all the electrons (six), then it will acquire the highest oxidation state +6 equal to the group number VI , if it takes the two necessary to complete the stable shell, it will acquire the lowest oxidation state –2 equal to Group number - 8 \u003d 6 - 8 \u003d -2.

3.3.3 Metal bond. Most metals have a number of properties that are general in nature and differ from the properties of other substances. Such properties are relatively high melting points, the ability to reflect light, high thermal and electrical conductivity. These features are explained by the existence in metals of a special type of interaction – metallic connection.

In accordance with the position in the periodic system, metal atoms have a small number of valence electrons, which are rather weakly bound to their nuclei and can easily be detached from them. As a result, positively charged ions appear in the crystal lattice of the metal, localized in certain positions of the crystal lattice, and a large number of delocalized (free) electrons that move relatively freely in the field of positive centers and carry out the connection between all metal atoms due to electrostatic attraction.

This is an important difference between metallic bonds and covalent bonds, which have a strict orientation in space. The bonding forces in metals are not localized and not directed, and the free electrons that form the "electron gas" cause high thermal and electrical conductivity. Therefore, in this case it is impossible to talk about the direction of the bonds, since the valence electrons are distributed almost uniformly over the crystal. This is precisely what explains, for example, the plasticity of metals, i.e., the possibility of displacement of ions and atoms in any direction

3.3.4 Donor-acceptor bond. In addition to the mechanism of formation of a covalent bond, according to which a common electron pair arises from the interaction of two electrons, there is also a special donor-acceptor mechanism . It lies in the fact that a covalent bond is formed as a result of the transition of an already existing (lone) electron pair donor (electron supplier) for the general use of the donor and acceptor (supplier of a free atomic orbital).

After formation, it is no different from covalent. The donor-acceptor mechanism is well illustrated by the formation of the ammonium ion (Figure 9) (asterisks indicate the electrons of the outer level of the nitrogen atom):

Figure 9 - Scheme of the formation of the ammonium ion

The electronic formula of the AVZ of the nitrogen atom is 2s 2 2p 3, that is, it has three unpaired electrons that enter into a covalent bond with three hydrogen atoms (1s 1), each of which has one valence electron. In this case, an ammonia molecule NH 3 is formed, in which the unshared electron pair of nitrogen is preserved. If a hydrogen proton (1s 0) that does not have electrons approaches this molecule, then nitrogen will transfer its pair of electrons (donor) to this hydrogen atomic orbital (acceptor), resulting in the formation of an ammonium ion. In it, each hydrogen atom is connected to the nitrogen atom by a common electron pair, one of which is realized by the donor-acceptor mechanism. It is important to note that the H-N bonds formed by various mechanisms do not have any differences in properties. This phenomenon is due to the fact that at the moment of bond formation, the orbitals of the 2s– and 2p– electrons of the nitrogen atom change their shape. As a result, four completely identical orbitals arise.

The donors are usually atoms with a large number of electrons, but with a small number of unpaired electrons. For elements of period II, in addition to the nitrogen atom, oxygen (two lone pairs) and fluorine (three lone pairs) have such a possibility. For example, the hydrogen ion H + in aqueous solutions is never in a free state, since the hydronium ion H 3 O + is always formed from water molecules H 2 O and the ion H +. The hydronium ion is present in all aqueous solutions, although for simplicity the spelling is preserved symbol H + .

3.3.5 Hydrogen bond. A hydrogen atom bonded to a strongly electronegative element (nitrogen, oxygen, fluorine, etc.), which “pulls” a common electron pair onto itself, experiences a shortage of electrons and acquires an effective positive charge. Therefore, it is able to interact with the lone pair of electrons of another electronegative atom (which acquires an effective negative charge) of the same (intramolecular bond) or another molecule (intermolecular bond). As a result, there is hydrogen bond , which is graphically indicated by dots:

This bond is much weaker than other chemical bonds (the energy of its formation is 10 – 40 kJ/mol) and mainly has a partly electrostatic, partly donor-acceptor character.

The hydrogen bond plays an extremely important role in biological macromolecules, such inorganic compounds as H 2 O, H 2 F 2, NH 3. For example, O-H bonds in H 2 O have a noticeable polar character with an excess of negative charge – on the oxygen atom. The hydrogen atom, on the contrary, acquires a small positive charge  + and can interact with lone pairs of electrons of the oxygen atom of the neighboring water molecule.

The interaction between water molecules turns out to be quite strong, such that even in water vapor there are dimers and trimers of the composition (H 2 O) 2, (H 2 O) 3, etc. In solutions, long chains of associates of this type can occur:

because the oxygen atom has two lone pairs of electrons.

The presence of hydrogen bonds explains the high boiling points of water, alcohols, carboxylic acids. Due to hydrogen bonds, water is characterized by such high melting and boiling points compared to H 2 E (E = S, Se, Te). If there were no hydrogen bonds, then water would melt at –100°C and boil at –80°C. Typical cases of association are observed for alcohols and organic acids.

Hydrogen bonds can occur both between different molecules and within a molecule if this molecule contains groups with donor and acceptor abilities. For example, it is intramolecular hydrogen bonds that play the main role in the formation of peptide chains that determine the structure of proteins. H-bonds affect the physical and chemical properties of a substance.

Hydrogen bonds do not form atoms of other elements , since the forces of electrostatic attraction of the opposite ends of the dipoles of polar bonds (О-Н, N-H, etc.) are rather weak and act only at small distances. Hydrogen, having the smallest atomic radius, allows such dipoles to approach each other so much that attractive forces become noticeable. No other element with a large atomic radius is capable of forming such bonds.

3.3.6 Forces of intermolecular interaction (van der Waals forces). In 1873, the Dutch scientist I. van der Waals suggested that there are forces that cause attraction between molecules. These forces were later called van der Waals forces. – the most versatile form of intermolecular bonding. The energy of the van der Waals bond is less than the hydrogen bond and is 2–20 kJ/∙mol.

Depending on the way the force is generated, they are divided into:

1) orientational (dipole-dipole or ion-dipole) - arise between polar molecules or between ions and polar molecules. When polar molecules approach each other, they are oriented in such a way that the positive side of one dipole is oriented towards the negative side of the other dipole (Figure 10).

Figure 10 - Orientation interaction

2) induction (dipole - induced dipole or ion - induced dipole) - arise between polar molecules or ions and non-polar molecules, but capable of polarization. Dipoles can act on non-polar molecules, turning them into indicated (induced) dipoles. (Figure 11).

Figure 11 - Inductive interaction

3) dispersive (induced dipole - induced dipole) - arise between non-polar molecules capable of polarization. In any molecule or atom of a noble gas, electric density fluctuations arise, as a result of which instantaneous dipoles appear, which in turn induce instantaneous dipoles in neighboring molecules. The movement of instantaneous dipoles becomes coordinated, their appearance and decay occur synchronously. As a result of the interaction of instantaneous dipoles, the energy of the system decreases (Figure 12).

Figure 12 - Dispersion interaction