Covalent bond length called the distance between the nuclei of atoms that form a bond. The bond length is directly related to the radius of the atom - the larger it is, the longer the bond.

The values ​​of the covalent radii of some atoms (pm; 10 -12 m):

• H = 30 pm;
• F=58;
• O = 73;
• N = 75;
• C=77;
• Cl = 99;
• S=103;
• P=110;
• Si = 118;
• Al = 130.

In symmetrical molecules (H 2, F 2, Cl 2 ...) half the bond length is called covalent radius. Knowing the covalent radius, it is very easy to calculate the length of a covalent bond in a molecule. For example, the length of the covalent bond of the HF molecule = 30 + 58 = 88 pm.

2. Energy of a covalent bond

Under covalent bond energy(expressed in kcal / mol or kJ / mol) usually understand the energy that is necessary to break the bond (when a covalent bond is formed, energy is released, when broken, it is absorbed). The higher the bond energy, the stronger the bond.

The bond energy depends on its length - the longer the bond in the molecule, the easier it is to break it (spend less energy).

Binding energies of some molecules (kJ/mol):

• H 2 = 453 (bond length = 60 pm);
• Cl 2 = 242 (198 pm);
• HCl = 431 (129 pm).

3. Polarity of a covalent bond

This characteristic displays the location of the electron pair of two atoms that form a bond. The degree of bond polarity depends on the magnitude of the electronegativity of the atoms forming the bond (the larger it is, the greater the bond polarity). In a more polar covalent bond, the shared pair of electrons is more biased toward the more electronegative atom (see the concept of electronegativity).

Electronegativity is a tabular value determined by the Polling scale. It is much more important to know not the electronegativity of the atom itself as such, but the difference between these values ​​in the molecule - which of the atoms is more electronegative, and which is less.

The polarity of a covalent bond is quantified using dipole moment(µ), while a system of two equivalent, but opposite in sign, charges is called dipole.

It is very important to distinguish between the dipole moment of a covalent bond (its polarity) and the dipole moment of the molecule as a whole. In simple diatomic molecules, these two parameters are equal to each other. A completely different picture is observed in complex molecules, in which the dipole moment of the molecule is the sum of the vectors of dipole moments of individual bonds.

4. Polarizability of a covalent bond

Polarizability refers to the degree to which electrons can be displaced by an external electric field generated by ions or other polar molecules.

The polarizability of a covalent bond is directly proportional to its length, which, in general, is logical - the farther an electron is from the nucleus of an atom, the weaker it is attracted to it, therefore, it is easier to shift under external influence on it. Thus, with an increase in the bond length, its polarizability increases, which, in turn, leads to an increase in the strength of acids (for example, hydroiodic acid is stronger than hydrofluoric acid).

The polarizability and polarity of a bond are inversely related: a less polar bond is more polarized, and vice versa.

5. Saturation of a covalent bond

Saturation is the ability of an atom to form a certain number of covalent bonds - all the "unpaired" electrons of the atom tend to take part in the formation of a bond. For example, a hydrogen atom has only one unpaired electron, while a nitrogen atom has three. For this reason, the most stable chemical compound will be NH 3, but not NH or NH 2.

6. Orientation of the covalent bond

Orientation characterizes the spatial orientation of a covalent bond relative to other bonds of the molecule. In molecules, the electrons of covalent bonds and free pairs of electrons constantly experience mutual repulsion, as a result of which the covalent bonds are located in such a way that the bond angle between them corresponds to the principle of least repulsion between electrons (for example, in a water molecule, the bond angle is 104.5 °).

7. Multiplicity of a covalent bond

In some cases, not one, but two (double bond) or three (triple bond) common electron pairs (the so-called multiple bonds) can occur between atoms.

A double covalent bond is formed in atoms that have two unpaired electrons; triple - for atoms that have three unpaired electrons (see Multiple bonds).

As can be seen from the table below, the nitrogen molecule is about 7 times "stronger" than the fluorine molecule.

Table of the dependence of the length and strength of a covalent bond on its multiplicity.

Communication length - internuclear distance. The shorter this distance, the stronger the chemical bond. The bond length depends on the radii of the atoms that form it: the smaller the atoms, the shorter the bond between them. For example, the length of the H-O bond is shorter than the length of the H-N bond (due to the smaller exchange of the oxygen atom).

An ionic bond is an extreme case of a polar covalent bond.

Metal connection.

The prerequisite for the formation of this type of connection is:

1) the presence of a relatively small number of electrons at the outer levels of atoms;

2) the presence of empty (vacant orbitals) at the outer levels of metal atoms

3) relatively low ionization energy.

Consider the formation of a metallic bond using sodium as an example. The valence electron of sodium, which is located on the 3s sublevel, can relatively easily move along the empty orbitals of the outer layer: along 3p and 3d. When atoms approach each other as a result of the formation of a crystal lattice, the valence orbitals of neighboring atoms overlap, due to which electrons move freely from one orbit to another, making a connection between ALL atoms of the metal crystal.

At the nodes of the crystal lattice there are positively charged ions and metal atoms, and between them there are electrons that can move freely throughout the crystal lattice. These electrons become common to all atoms and ions of the metal and are called "electron gas". The bond between all positively charged metal ions and free electrons in the crystal lattice of metals is called metallic bond.

The presence of a metallic bond determines the physical properties of metals and alloys: hardness, electrical conductivity, thermal conductivity, malleability, ductility, metallic luster. Free electrons can carry heat and electricity, so they are the cause of the main physical properties that distinguish metals from non-metals - high electrical and thermal conductivity.

Hydrogen bond.

hydrogen bond occurs between molecules that include hydrogen and atoms with high EO (oxygen, fluorine, nitrogen). The covalent bonds H-O, H-F, H-N are strongly polar, due to which an excess positive charge accumulates on the hydrogen atom, and an excess negative charge accumulates on the opposite poles. Electrostatic attraction forces arise between oppositely charged poles - hydrogen bonds.

Hydrogen bonds can be both intermolecular and intramolecular. The energy of a hydrogen bond is about ten times less than the energy of a conventional covalent bond, but nevertheless, hydrogen bonds play an important role in many physicochemical and biological processes. In particular, DNA molecules are double helixes in which two chains of nucleotides are linked by hydrogen bonds. Intermolecular hydrogen bonds between water and hydrogen fluoride molecules can be depicted (dots) as follows:

Substances with a hydrogen bond have molecular crystal lattices. The presence of a hydrogen bond leads to the formation of associates of molecules and, as a consequence, to an increase in the melting and boiling points.

In addition to the listed main types of chemical bonds, there are also universal forces of interaction between any molecules that do not lead to the breaking or formation of new chemical bonds. These interactions are called van der Waals forces. They cause the attraction of the molecules of a given substance (or various substances) to each other in liquid and solid states of aggregation.

Different types of chemical bonds determine the existence of different types of crystal lattices (table).

Molecular substances have molecular structure. Such substances include all gases, liquids, as well as solid substances with a molecular crystal lattice, such as iodine. Solids with an atomic, ionic, or metallic lattice have non-molecular structure, they do not contain molecules.

Table

Feature of the crystal lattice	Crystal lattice type
Molecular	Ionic	Atomic	metal
Particles at lattice sites	molecules	Cations and anions	atoms	Cations and metal atoms
The nature of the connection between particles	Forces of intermolecular interaction (including hydrogen bonds)	Ionic bonds	covalent bonds	metal connection
Bond strength	Weak	durable	Very durable	different strength
Distinctive physical properties of substances	Fusible or subliming, low hardness, many soluble in water	Refractory, hard, brittle, many soluble in water. Solutions and melts conduct electricity	Very refractory, very hard, practically insoluble in water	High electrical and thermal conductivity, metallic luster, ductility.
Substance examples	Simple substances - non-metals (in the solid state): Cl 2, F 2, Br 2, O 2, O 3, P 4, sulfur, iodine, (except silicon, diamond, graphite); complex substances consisting of non-metal atoms (except ammonium salts): water, dry ice, acids, non-metal halides: PCl 3, SiF 4, CBr 4, SF 6, organic substances: hydrocarbons, alcohols, phenols, aldehydes, etc.	Salts: sodium chloride, barium nitrate, etc.; alkalis: potassium hydroxide, calcium hydroxide, ammonium salts: NH 4 Cl, NH 4 NO 3, etc., metal oxides, nitrides, hydrides, etc. (compounds of metals with non-metals)	Diamond, graphite, silicon, boron, germanium, silicon oxide (IV) - silica, SiC (carborundum), black phosphorus (P).	Copper, potassium, zinc, iron and other metals
Comparison of substances by melting and boiling points.
	Due to the weak forces of intermolecular interaction, such substances have the lowest melting and boiling points. Moreover, the greater the molecular weight of the substance, the higher t 0 pl. it has. Exceptions are substances between the molecules of which hydrogen bonds can form. For example, HF has a higher t 0 pl. than HCl.	Substances have high t 0 pl., but lower than substances with an atomic lattice. The higher the charges of the ions that are at the lattice sites and the shorter the distance between them, the higher the melting point of the substance. For example, t 0 square. CaF 2 is higher than t 0 pl. KF.	They have the highest t 0 pl. The stronger the bond between the atoms in the lattice, the higher t 0 pl. has substance. For example, Si has a lower t 0 square than C.	Metals have different t0 pl.: from -37 0 С for mercury to 3360 0 С for tungsten.

Definition

A covalent bond is a chemical bond formed due to the socialization of atoms of their valence electrons. An obligatory condition for the formation of a covalent bond is the overlap of atomic orbitals (AO), on which valence electrons are located. In the simplest case, the overlap of two AOs leads to the formation of two molecular orbitals (MOs): a bonding MO and an antibonding (loosening) MO. Shared electrons are located on a lower energy binding MO:

Communication education

Covalent bond (atomic bond, homeopolar bond) - a bond between two atoms due to the socialization (electron sharing) of two electrons - one from each atom:

A. + B. -> A: B

For this reason, the homeopolar relationship has a directional character. A pair of electrons making a bond belongs simultaneously to both bonding atoms, for example:

	..		..						..
:	Cl	:	Cl	:			H	:	O	:	H
	..		..						..

Types of covalent bond

There are three types of covalent chemical bonds that differ in the mechanism of their formation:

1. Simple covalent bond. For its formation, each of the atoms provides one unpaired electron. When a simple covalent bond is formed, the formal charges of the atoms remain unchanged. If the atoms forming a simple covalent bond are the same, then the true charges of the atoms in the molecule are also the same, since the atoms forming the bond equally own a socialized electron pair, such a bond is called a non-polar covalent bond. If the atoms are different, then the degree of ownership of a socialized pair of electrons is determined by the difference in the electronegativity of the atoms, an atom with a greater electronegativity has a pair of bond electrons to a greater extent, and therefore its true charge has a negative sign, an atom with a lower electronegativity acquires, respectively, the same charge, but with a positive sign.

Sigma (σ)-, pi (π)-bonds - an approximate description of the types of covalent bonds in the molecules of organic compounds, σ-bond is characterized by the fact that the density of the electron cloud is maximum along the axis connecting the nuclei of atoms. When a π-bond is formed, the so-called lateral overlap of electron clouds occurs, and the density of the electron cloud is maximum "above" and "below" the plane of the σ-bond. For example, take ethylene, acetylene and benzene.

In the ethylene molecule C 2 H 4 there is a double bond CH 2 \u003d CH 2, its electronic formula is: H: C:: C: H. The nuclei of all ethylene atoms are located in the same plane. Three electron clouds of each carbon atom form three covalent bonds with other atoms in the same plane (with angles between them of about 120°). The cloud of the fourth valence electron of the carbon atom is located above and below the plane of the molecule. Such electron clouds of both carbon atoms, partially overlapping above and below the plane of the molecule, form a second bond between carbon atoms. The first, stronger covalent bond between carbon atoms is called a σ-bond; the second, less strong covalent bond is called a π-bond.

In a linear acetylene molecule

H-S≡S-N (N: S::: S: N)

there are σ-bonds between carbon and hydrogen atoms, one σ-bond between two carbon atoms, and two π-bonds between the same carbon atoms. Two π-bonds are located above the sphere of action of the σ-bond in two mutually perpendicular planes.

All six carbon atoms of the C 6 H 6 cyclic benzene molecule lie in the same plane. σ-bonds act between carbon atoms in the plane of the ring; the same bonds exist for each carbon atom with hydrogen atoms. Each carbon atom spends three electrons to make these bonds. Clouds of the fourth valence electrons of carbon atoms, having the shape of eights, are located perpendicular to the plane of the benzene molecule. Each such cloud overlaps equally with the electron clouds of neighboring carbon atoms. In the benzene molecule, not three separate π-bonds are formed, but a single π-electron system of six electrons, common to all carbon atoms. The bonds between the carbon atoms in the benzene molecule are exactly the same.

A covalent bond is formed as a result of the socialization of electrons (with the formation of common electron pairs), which occurs during the overlap of electron clouds. Electron clouds of two atoms participate in the formation of a covalent bond. There are two main types of covalent bonds:

• A covalent non-polar bond is formed between non-metal atoms of the same chemical element. Simple substances have such a bond, for example, O 2; N 2 ; C 12 .
• A covalent polar bond is formed between atoms of different non-metals.

see also

Literature

• "Chemical Encyclopedic Dictionary", M., "Soviet Encyclopedia", 1983, p.264.

Organic chemistry
List of organic compounds

Structural chemistry
Chemical bond:	Aromaticity | covalent bond| Ionic bond | Metal connection | Hydrogen bond | Donor-acceptor bond | Tautomerism
Structure display:	Functional group | Structural formula | Chemical formula | ligand
Electronic properties:	Electronegativity | Electron affinity | Ionization energy | Dipole | Octet rule
Stereochemistry:	Asymmetric atom | Isomerism | Configuration | Chirality | Conformation

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• Great Polytechnic Encyclopedia

CHEMICAL BOND The mechanism by which atoms combine to form molecules. There are several types of such a bond, based either on the attraction of opposite charges, or on the formation of stable configurations through the exchange of electrons. ... ... Scientific and technical encyclopedic dictionary

chemical bond- CHEMICAL BOND, the interaction of atoms, causing their connection into molecules and crystals. The forces acting during the formation of a chemical bond are mainly electrical in nature. The formation of a chemical bond is accompanied by a rearrangement ... ... Illustrated Encyclopedic Dictionary

Mutual attraction of atoms, leading to the formation of molecules and crystals. It is customary to say that in a molecule or in a crystal between neighboring atoms there are ch. The valence of an atom (which is discussed in more detail below) indicates the number of bonds ... Great Soviet Encyclopedia

chemical bond- mutual attraction of atoms, leading to the formation of molecules and crystals. The valence of an atom shows the number of bonds formed by a given atom with neighboring ones. The term "chemical structure" was introduced by Academician A. M. Butlerov in ... ... Encyclopedic Dictionary of Metallurgy

An ionic bond is a strong chemical bond formed between atoms with a large difference in electronegativity, in which a common electron pair is completely transferred to an atom with a greater electronegativity. An example is the CsF compound ... Wikipedia

Chemical bonding is a phenomenon of the interaction of atoms, due to the overlap of electron clouds, binding particles, which is accompanied by a decrease in the total energy of the system. The term "chemical structure" was first introduced by A. M. Butlerov in 1861 ... ... Wikipedia

The covalent bond is characterized orientation in space, polarity, multiplicity, energy and length.

As we know, electron orbitals (except s-orbitals) have spatial orientation. The covalent bond, which is the result of electron-nuclear interactions, is located in a certain direction with respect to the nuclei of these atoms. If electron clouds overlap in the direction of the straight line that connects the nuclei of atoms (i.e., along the bond axis), such a covalent bond is called s-bond(sigma bond). For example, in H 2 , Cl 2 , HC1 molecules, the atoms are connected by a covalent s-bond. Covalent sigma bonds are formed when orbitals overlap: s- s (as in H 2): s - R(as in HC1), R- R(as in C1 2).

When the p-orbitals directed perpendicular to the bond axis overlap, two overlapping regions are formed on both sides of the bond axis. Such a covalent bond is called a p-bond (pi-bond) (Fig. 6). For example, in a nitrogen molecule, the atoms are linked by one s-bond and two p-bonds (Fig. 7).

Rice. 6. Schematic representation of the p-bond

Rice. 7. Schematic representation of s- and p-bonds in a nitrogen molecule

The orientation of the covalent bond determines the spatial structure of the molecules, i.e., their shape. The hydrogen chloride molecule has a linear shape: it is formed using one s-bond (s - p-orbitals). The water molecule has an angular structure: it is formed due to the overlap of the s-orbitals of two hydrogen atoms with two mutually perpendicular p-orbitals of the oxygen atom (Fig. 8). Therefore, the angle between s-bonds in a water molecule must be equal to 90°. In fact, the angle is 104.5°, which is explained by the phenomenon of hybridization. The ammonia molecule has the shape of a regular pyramid, the methane molecule has the shape of a tetrahedron.

Rice. 8. The structure of the water molecule

Communication polarity is determined by the asymmetry in the distribution of the common electron cloud along the bond axis.

If common electron pairs are located symmetrically with respect to both nuclei, then such a covalent bond is called non-polar.

In the molecules of simple substances - hydrogen H 2, oxygen O 2, nitrogen N 2, chlorine C1 2, fluorine F 2, atoms are connected by a non-polar covalent bond.

If common electron pairs are shifted to one of the atoms (they are located asymmetrically with respect to the nuclei of various atoms), then such a covalent bond is called polar.

The bond in the molecules of water H 2 O, ammonia NH 3, hydrogen chloride HC1 is polar.

multiplicity covalent bond is determined by the number of shared electron pairs that link atoms.

The bond between two atoms using one pair of electrons is called simple(bonds H - C1, C - H, H - O, etc.). A bond between two atoms using two electron pairs is called double. A bond between two atoms using three electron pairs is called triple.

For example, a double bond is observed between carbon atoms in ethylene H 2 C \u003d CH 2, a triple bond is observed in nitrogen molecules N N, acetylene H - C C - H.

Link length is the equilibrium distance between the nuclei of atoms. The bond length is expressed in nanometers (nm). The shorter the bond length, the stronger the chemical bond. The strength of a bond is measured by its energy.

Bond energy is equal to the work that must be expended to break the connection. Express the binding energy in kilojoules per mol (kJ/mol); for example, in a hydrogen molecule, the binding energy is 435 kJ/mol. The bond energy increases with decreasing bond length (Table 10).

Table 10 Type, length and energy of bonds in the molecules of certain substances

The bond energy increases with increasing bond multiplicity (Table 11).

Table 11 Bond length and energy between nitrogen atoms and between carbon atoms

The process of bond formation proceeds with the release of energy (exothermic process), and the process of breaking the bond - with the absorption of energy (endothermic process).

Polarity of molecules

Polarity of molecules depends on the polarity of the individual bonds and on their location in the molecule (i.e., on the structure of the molecules).

Molecules of simple substances (H 2, F 2, N 2, etc.) formed by non-polar covalent bonds, non-polar.

Molecules of complex substances can be both non-polar and polar. Examples of substances with non-polar molecules: carbon dioxide CO 2, methane CH 4, benzene C 6 H 6, glucose C 6 H 12 O 6, dimethyl ether C 2 H 6 O, etc. Examples of substances with polar molecules: sulfur dioxide SO 2, water H 2 O, ammonia NH 3, ethyl alcohol C 2 H 5 OH, etc.

In non-polar molecules, the "center of gravity" of the electron cloud coincides with the "center of gravity" of the positive charge of the nuclei. In polar molecules, the "center of gravity" of the electron cloud does not coincide with the "center of gravity" of the positive charge.

For example, in a HC1 hydrogen chloride molecule, the electron density near the chlorine nucleus is higher than near the hydrogen nucleus, i.e., the chlorine atom has a negative charge q = - 0.18, and the hydrogen atom has a positive charge q-= + 0.18. Charges (q) atoms in a molecule are called .efficient. Therefore, polar molecules can be considered as electric dipoles, in which charges, different in sign but equal in magnitude, are located at a certain distance from each other. The measure of polarity of molecules is electric moment of the dipole.

The electric moment of a dipole is the product of the effective charge times the distance between the centers of positive and negative charges in the molecule. The electric moment of a dipole in a molecule depends on its structure. The presence or absence of the electric moment of the dipole makes it possible to judge the geometric structure of the molecule. For example, the CO 2 molecule is non-polar, while the SO 2 molecule has an electric dipole moment. It follows that the CO 2 molecule has a linear structure, and the SO 2 molecule has an angular structure.

The properties of substances depend on the polarity of the molecules. Substances whose molecules are polar have higher boiling and melting points than substances whose molecules are nonpolar. This is due to the mutual attraction of polar molecules.

Electronegativity

The ability of atoms of a chemical element to attract common electron pairs is called electronegativity.

The electronegativity of an element is determined by the sum of its ionization energy and electron affinity. The relative electronegativity of the atoms of some elements are given in Table. 12.

Table 12 Relative electronegativity of some elements

Period	Group
	I	II	III	IV	V	VI	VII
	H 2.1
	Li 0.98	Be 1.5	In 2.0	From 2.5	N 3.07	About 3.50	F4.0
	Na 0.93	Mg 1.2	Al 1.6	Si 1.9	P 2.2	S 2.6	Cl 3.0
	K 0.91	Ca 1.04	Ga 1.8	Ge 2.0	As 2.1	Se 2.5	Br2.8
	Rb 0.89	Sr 0.99	In 1.5	sn 1.7	Sb 1.8	Those 2.1	I 2.6

The greater the electronegativity of an atom, the stronger it attracts a common electron pair. When a covalent bond is formed between two atoms of different elements, the shared electron pairs shift towards the more electronegative atom. For example, in the water molecule H 2 O, the common electron pairs are shifted to the oxygen atom.

The relative electronegativity of an atom is not a strictly constant value and is used only to determine the direction of displacement of common electron pairs during the formation of molecules

The electronegativity of elements obeys the periodic law. In a period, the electronegativity of elements increases with the increase in the atomic number of the element. At the beginning of the period, there are elements with low electronegativity (metals), and at the end of the period, elements with the highest electronegativity (non-metals). In the subgroup, the electronegativity of elements decreases with increasing serial number. The most electronegative element in the periodic table is fluorine. Inert elements have no electronegativity.

Chemical elements can be arranged in a row in ascending order of electronegativity.

Sb, Si. B, As. H, Te. R. C, Se, I, S, Br. Cl, N. O, F

electronegativity increases

Electronegativity characterizes the difference in the properties of elements. Therefore, it is used as a qualitative characteristic in determining the nature of a chemical bond in various compounds.

Ionic bond

When compounds are formed from elements that are very different in electronegativity (typical metals and typical non-metals), the common electron pairs are completely shifted to the more electronegative atom. As a result, ions.

For example, during the combustion of sodium in chlorine, the unpaired 3s electron of the sodium atom pairs with the 3p electron of the chlorine atom. The shared electron pair shifts completely to the chlorine atom. As a result, sodium ion Na + and chloride ion CI - are formed.

Charged particles, into which atoms turn as a result of the return or addition of electrons, are called ions.

The charge of a negative ion is equal to the number of electrons that the atom has attached. The charge of a positive ion is equal to the number of electrons that the atom has donated.

Oppositely charged nonons attract each other.

Compounds that are formed from ions are called ionic. The bond between ions is called ionic.

There is no sharp boundary between ionic and covalent bonds. An ionic bond can be considered as an extreme case of a covalent polar bond (Fig. 9). Unlike covalent bonds, ionic bonds are non-directional.

The process of donating electrons is called oxidation. The process of adding electrons is called reduction.

For example, when sodium reacts with chlorine, the sodium atom donates an electron, is oxidized and sodium ion Na is formed - e-®Na +

Figure 9. Schematic of the transition from a covalent bond to a kyonic

The chlorine atom attaches an electron, is reduced and the chloride ion Cl + is formed e -®Cl - .

Metals of the main subgroups of groups I and II, when combined with non-metals of the main subgroup of group VII, form typical ionic compounds. For example, sodium chloride NaCl, potassium fluoride KF, calcium chloride CaCl 2.

Ionic compounds are solid crystalline substances.

hydrogen bond

A hydrogen atom bonded to a highly electronegative element (fluorine, oxygen, nitrogen) can form another bond to another atom of a highly electronegative element. For example, in a water molecule, hydrogen atoms are linked to oxygen atoms by a polar covalent bond. Shared electron pairs are displaced towards the oxygen atom. The hydrogen atom has a partial positive charge, and the oxygen atom has a partial negative charge. The positively charged hydrogen atom of one water molecule is attracted to the negatively charged oxygen atom of another water molecule. Between two oxygen atoms there is a bond formed with the help of a hydrogen atom. The hydrogen atom is on a straight line that connects the nuclei of these atoms

O ¾ H. . . O ¾ H. . . O ¾ H. . . O ¾ H

A hydrogen bond is formed due to the forces of electrostatic attraction of polar molecules to each other, especially when they contain atoms of strongly electronegative elements (F, O, N).

For example, hydrogen bonds form HF, H 2 O, NH 3, but do not form their analogs HCl, H 2 S, PH 3.

Hydrogen bonds are unstable and break quite easily (for example, when ice melts and water boils), but since some energy is required to break these bonds, the melting and boiling points of substances with hydrogen bonds between molecules turn out to be much higher than those of similar substances, but without hydrogen bonds. For example:

(in HF and H 2 O there are hydrogen bonds, but in HCl and H 2 S they are not).

Many organic compounds also form hydrogen bonds, and hydrogen bonds play an important role in biological processes.

metal connection

Metals have the lowest ionization energy. Therefore, in metals, valence electrons are easily detached from individual atoms and become common to the entire crystal. (socialised). This is how positive metal ions are formed and electron gas- set of mobile electrons. In a metal crystal, a small number of shared electrons bind a large number of ions.

The chemical bond in metals between positive ions and socialized electrons is called a metallic bond.

A metallic bond is similar to a covalent bond. The formation of these bonds is based on the processes of socialization of valence electrons. But in a metal, valence electrons are common to the entire crystal, and in compounds with a covalent bond, only valence electrons of two neighboring atoms are common. The metallic bond is non-directional, since the valence electrons are distributed almost evenly throughout the crystal.

A metallic bond is characteristic only of metals in a solid or liquid state of aggregation.

SOLUTIONS

Similar information.

Introduction. 3

1 Covalent bond. Basic concepts. 4

2 The main characteristics of the covalent bond. 6

3 Types of covalent bond. eight

4 Valency. ten

Introduction

A relatively small number of elements in the periodic system of Dmitri Ivanovich Mendeleev - 118 - form about 10 million simple and complex substances. The reason for this phenomenon lies in the fact that, interacting with each other, the atoms of many elements bind to each other, forming different chemical compounds.

The force that connects two or more interacting atoms into molecules or other particles is called a chemical bond.

The reason for the formation of a chemical bond is the desire of metal and non-metal atoms to achieve a more stable electronic structure by interacting with other atoms. When a chemical bond is formed, the electronic structures of the bonding atoms are significantly rearranged, therefore, many of their properties in compounds change.

In the word "covalent" the prefix "co-" means "joint participation". And "valenta" in translation into Russian - strength, ability. In this case, we mean the ability of atoms to bond with other atoms. One example of a chemical bond is a covalent bond.

The term covalent bond was first coined by Nobel laureate Irving Langmuir in 1919. The term referred to a chemical bond due to the shared possession of electrons, as opposed to a metallic bond, in which the electrons were free, or an ionic bond, in which one of the atoms donated an electron and became a cation, and the other atom accepted an electron and became an anion.

Later (1927), F. London and W. Heitler, using the example of a hydrogen molecule, gave the first description of a covalent bond from the point of view of quantum mechanics.

covalent bond. Basic concepts

When a covalent bond is formed, atoms combine their electrons, as it were, into a common "piggy bank" - a molecular orbital, which is formed from the atomic shells of individual atoms. This new shell contains as many complete electrons as possible and replaces the atoms with their own incomplete atomic shells.

Let us consider the emergence of a covalent bond using the example of the formation of a hydrogen molecule from two hydrogen atoms (Fig. 1). This process is already a typical chemical reaction, because from one substance (atomic hydrogen) another is formed - molecular hydrogen. An external sign of the energy efficiency of this process is the release of a large amount of heat.

Rice. 1. The emergence of a covalent bond during the formation of a hydrogen molecule from two hydrogen atoms.

The electron shells of hydrogen atoms (with one s-electron for each atom) merge into a common electron cloud (molecular orbital), where both electrons "serve" the nuclei, regardless of whether this nucleus is "own" or "foreign".

When the electron shells of two hydrogen atoms approach and form a new, now molecular electron shell (Fig. 1), this new shell is similar to the completed electron shell of the noble gas atom helium.

Completed shells, as we remember, are more stable than unfinished ones. Thus, the total energy of the new system, the hydrogen molecule, turns out to be much lower than the total energy of two unbound hydrogen atoms. The excess energy is released in the form of heat.

In the resulting system of two hydrogen atoms, each nucleus is served by two electrons. In a new (molecular) shell, it is no longer possible to distinguish which of the electrons previously belonged to one or another atom. It is customary to say that electrons are socialized. Since both nuclei claim a pair of electrons equally, the electron density is concentrated both around the nuclei and in the space between the atoms (this is shown in Fig. 2).

Rice. 2. Another way to depict atomic and molecular orbitals

In Figure 2, the density of dots reflects the "electron density", that is, the probability of finding an electron at any point in space near the nuclei of hydrogen atoms. It can be seen that a significant electron density is concentrated in the space between the two nuclei in the hydrogen molecule.

A covalent bond is the bonding of atoms with the help of common (shared between them) electron pairs. 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.

Main characteristics of a covalent bond

The main characteristics of a covalent bond are: bond length (distance between the centers of atoms in a molecule); bond energy (the energy that must be spent to break the bond); bond polarity (uneven distribution of electron density between atoms due to different electronegativity); polarizability (the ease with which the electron density of the bond to one of the atoms is swept away under the influence of external factors); orientation (covalent bond directed to the line connecting the centers of atoms).

The direction of the bond is due to the molecular structure of the substance and the geometric shape of their molecule. The angles between two bonds are called bond angles.

Saturation - the ability of atoms to form a limited number of covalent bonds. The number of bonds formed by an atom is limited by the number of its outer atomic orbitals.

The polarity of the bond is due to the uneven distribution of the electron density due to differences in the electronegativity of the atoms. On this basis, covalent bonds are divided into non-polar and polar.

The polarizability of a bond is expressed in the displacement of bond electrons under the influence of an external electric field, including that of another reacting particle. Polarizability is determined by the electron mobility. The polarity and polarizability of covalent bonds determine the reactivity of molecules with respect to polar reagents. Electrons are more mobile the farther they are from nuclei.

Depending on the electronegativity of the atoms between which a covalent bond has formed, it can be polar or non-polar.

If the electronegativity of the atoms is the same, then the common electron pair is at the same distance from the nucleus of each of the atoms. Such a bond is called covalent-nonpolar. When a covalent bond occurs between atoms with different electronegativity, the common electron pair is shifted to a more electronegative atom. In this case, a covalent polar bond is formed. The arrow in the formula indicates the polarity of the covalent bond. Using the Greek letter b (“delta”), partial charges on atoms are denoted: b + - reduced, 6 - increased electron density.

According to the number of electron pairs forming a covalent bond, simple bonds are distinguished - with one pair of electrons and multiple - with two or three pairs.