(Based on materials from the site http://chemel.ru/2008-05-24-19-19-34/2008-06-01-15-23-43/18-2008-05-29-22-08-32. html)
It is known that non-metals interact with each other. Consider the mechanism of the emergence of a covalent bond using the example of the formation of a hydrogen molecule:
H + H \u003d H 2 H \u003d - 436kJ / mol
Imagine that we have two separate isolated hydrogen atoms. The nucleus of each of the free hydrogen atoms is surrounded by a spherical symmetric electron cloud formed by a 1s electron (see Fig. 1). When the atoms approach each other up to a certain distance, the electron shells (orbitals) partially overlap (Fig. 2).
As a result, a molecular two-electron cloud arises between the centers of both nuclei, which has a maximum electron density in the space between the nuclei; an increase in the density of the negative charge favors a strong increase in the forces of attraction between the nuclei and the molecular cloud.
So, a covalent bond is formed as a result of the overlapping of electron clouds of atoms, accompanied by the release of energy. If for hydrogen atoms approaching before touching, the distance between the nuclei is 0.106 nm, then after the overlap of electron clouds (formation of the H 2 molecule), this distance is 0.074 nm (Fig. 2).
Usually, the greatest overlap of electron clouds occurs along the line connecting the nuclei of two atoms.
The stronger the chemical bond, the greater the overlap of electron orbitals.
As a result of the occurrence of a chemical bond between two hydrogen atoms, each of them reaches the electronic configuration of a noble gas atom.
Depicting chemical bonds is customary in different ways:
1) with the help of electrons in the form of dots placed at the chemical sign of the element.
Then the formation of a hydrogen molecule can be shown by the scheme:
H + H H:H
2) with the help of quantum cells (Hund cells), as the placement of two electrons with opposite spins in one molecular quantum cell:
The diagram on the left shows that the molecular energy level is lower than the original atomic levels, which means that the molecular state of a substance is more stable than the atomic state.
3) often, especially in organic chemistry, a covalent bond is represented by a dash (dash)
(for example, H-H), which symbolizes a pair of electrons.
A covalent bond in a chlorine molecule is also carried out using two common electrons, or an electron pair: 
As you can see, each chlorine atom has three lone pairs and one unpaired electron.
The formation of a chemical bond occurs due to the unpaired electrons of each atom. Unpaired electrons are bound into a common pair of electrons, also called a common (shared) pair.
If one covalent bond has arisen between atoms (one common electron pair), then it is called single; if more, then multiple (two common electron pairs), triple (three common electron pairs).
A single bond is represented by one dash (stroke), a double bond by two, and a triple bond by three. A dash between two atoms shows that they have a pair of electrons generalized, as a result of which a chemical bond was formed. With the help of such dashes, the sequence of connection of atoms in a molecule is depicted.
So, in the chlorine molecule, each of its atoms has a completed external level of eight electrons (s 2 p 6), and two of them (an electron pair) equally belong to both atoms.
The bond in the oxygen molecule O 2 is depicted somewhat differently. It has been experimentally established that oxygen is a paramagnetic substance (it is drawn into a magnetic field). Its molecule has two unpaired electrons. The structure of this molecule can be represented as follows:
An unambiguous solution to the image of the electronic structure of the oxygen molecule has not yet been found. However, it cannot be shown like this:
In the nitrogen molecule N 2, atoms have three common electron pairs:
It is obvious that the nitrogen molecule is stronger than the oxygen or chlorine molecule, which is the reason for the significant inertness of nitrogen in chemical reactions.
A chemical bond carried out by electron pairs is called a covalent bond.
It is a two-electron and two-center (holds two nuclei) bond.
Compounds with a covalent bond are called homeopolar, or atomic.
There are two types of covalent bonds: non-polar and polar.
In the case of a non-polar covalent bond, the electron cloud formed by a common pair of electrons, or the bond electron cloud, is distributed in space symmetrically with respect to the nuclei of both atoms.
An example is diatomic molecules consisting of atoms of one element: H 2 Cl 2, O 2, N 2, F 2, etc., in which the electron pair equally belongs to both atoms.
In the case of a polar covalent bond, the electron cloud of the bond is shifted towards the atom with a higher relative electronegativity.
Molecules of volatile inorganic compounds can serve as an example: HC1, H 2 O, H 2 S, NH 3, etc.
The formation of the HC1 molecule can be represented by the scheme:
The electron pair is shifted to the chlorine atom, since the relative electronegativity of the chlorine atom (2.83) is greater than that of the hydrogen atom (2.1).
A covalent bond is formed not only due to the overlap of one-electron clouds, it is an exchange mechanism for the formation of a covalent bond.
Another mechanism for the formation of a covalent bond is also possible - a donor-acceptor one. In this case, the chemical bond arises due to the two-electron cloud of one atom and the free orbital of another atom. Let us consider as an example the mechanism of formation of the ammonium ion NH +4 . In the ammonia molecule, the nitrogen atom has an unshared pair of electrons (two-electron-
cloud):
The hydrogen ion has a free (not filled) 1s-orbital, which can be denoted as follows: H +. When an ammonium ion is formed, a two-electron cloud of nitrogen becomes common for nitrogen and hydrogen atoms, i.e. it turns into a molecular electron cloud. So, there is a fourth covalent bond.
The process of formation of the ammonium ion can be represented by the scheme:
The charge of the hydrogen ion becomes common (it is delocalized, i.e. dispersed between all atoms), and the two-electron cloud (lone electron pair) belonging to nitrogen becomes common with hydrogen. In diagrams, the cell image is often omitted.
An atom that donates a lone electron pair is called a donor, and an atom that accepts it (i.e., provides a free orbital) is called an acceptor.
The mechanism of formation of a covalent bond due to a two-electron cloud of one atom (donor) and a free orbital of another atom (acceptor) is called donor-acceptor. The covalent bond formed in this way is called a donor-acceptor, or coordination, bond.
However, this is not a special type of bond, but only a different mechanism (method) for the formation of a covalent bond. The properties of the fourth N-H-bond in the ammonium ion are no different from the rest of the bonds.
metal connection
The atoms of most metals at the outer energy level contain a small number of electrons. So, one electron each contains 16 elements, two - 58, three - 4 elements, and none - only in Pd. The atoms of the elements Ge, Sn and Pb have 4 electrons at the outer level, Sb and Bi - 5 each, Po - 6, but these elements are not characteristic metals.
Metal elements form simple substances - metals. Under normal conditions, these are crystalline substances (except mercury). On fig. 3 is a diagram of the crystal lattice of sodium.
As you can see, each sodium atom is surrounded by eight neighboring ones. Using sodium as an example, consider the nature of the chemical bond in metals.
The sodium atom, like other metals, has an excess of valence orbitals and a lack of electrons.
So, a valence electron (3s 1) can occupy one of nine free orbitals - 3s (one), Zp (three) and 3d (five).
When approaching atoms as a result the crystal lattice, the valence orbitals of neighboring atoms overlap,
due to which electrons move freely from one orbital to another, making a connection between all atoms of the metal crystal. This type of chemical bond is called a metallic bond.
A metallic bond is formed by elements whose atoms at the outer level have few valence electrons compared to the total number of outer energetically close orbitals, and valence electrons are weakly retained in the atom due to the low ionization energy.
The chemical bond in metal crystals is highly delocalized, i.e. the electrons that carry out the connection are socialized (“electron gas”) and move around the whole piece of metal, which is generally electrically neutral.
The metallic bond is characteristic of metals in the solid and liquid state. This is a property of aggregates of atoms located in close proximity to each other. However, in the vapor state, metal atoms, like all substances, are linked by a covalent bond. Metal pairs consist of individual molecules (monatomic and diatomic). The bond strength in a crystal is greater than in a metal molecule, and therefore the process of formation of a metal crystal proceeds with the release of energy.
A metallic bond bears some resemblance to a covalent bond, since it is also based on the socialization of valence electrons. However, the electrons that carry out the covalent bond are located near the connected atoms and are strongly associated with them. The electrons that carry out the metallic bond move freely throughout the crystal and belong to all its atoms. That is why crystals with a covalent bond are brittle, and those with a metal bond are plastic, i.e. they change shape on impact, are rolled into thin sheets, and drawn into wire.
The metallic bond explains the physical properties of metals.
hydrogen bond
A hydrogen bond is a kind of chemical bond. It can be intermolecular and intramolecular.
An intermolecular hydrogen bond occurs between molecules that include hydrogen and a strongly electronegative element - fluorine, oxygen, nitrogen, less often chlorine, sulfur. Since in such a molecule the common electron pair is strongly shifted from hydrogen to the atom of the electronegative element, and the positive charge of hydrogen is concentrated in a small volume, the proton interacts with the unshared electron pair of another atom or ion, socializing it. As a result, a second, weaker bond is formed, called a hydrogen bond.
Previously, the hydrogen bond was reduced to an electrostatic attraction between a proton and another polar group. But it should be considered more correct that the donor-acceptor interaction also contributes to its formation. This connection is characterized by orientation in space and saturation.
Usually, a hydrogen bond is indicated by dots and this indicates that it is much weaker than a covalent bond (about 15-20 times). Nevertheless, it is responsible for the association of molecules. For example, the formation of dimers (they are most stable in the liquid state) of water and acetic acid can be represented by the following schemes:
As can be seen from these examples, two water molecules are combined by hydrogen bonding, and in the case of acetic acid, two acid molecules are combined to form a cyclic structure.
The presence of hydrogen bonds explains the higher boiling point of water (100 ° C) compared with hydrogen compounds of elements of the oxygen subgroup ( H 2 O, H 2 S, H 2 Te). In the case of water, additional energy must be expended to break hydrogen bonds.
It is extremely rare for chemical substances to consist of individual, unrelated atoms of chemical elements. Under normal conditions, only a small number of gases called noble gases have such a structure: helium, neon, argon, krypton, xenon and radon. Most often, chemical substances do not consist of disparate atoms, but of their combinations into various groups. Such combinations of atoms can include several units, hundreds, thousands, or even more atoms. The force that keeps these atoms in such groupings is called chemical bond.
In other words, we can say that a chemical bond is an interaction that ensures the bonding of individual atoms into more complex structures (molecules, ions, radicals, crystals, etc.).
The reason for the formation of a chemical bond is that the energy of more complex structures is less than the total energy of the individual atoms that form it.
So, in particular, if an XY molecule is formed during the interaction of X and Y atoms, this means that the internal energy of the molecules of this substance is lower than the internal energy of the individual atoms from which it was formed:
E(XY)< E(X) + E(Y)
For this reason, when chemical bonds are formed between individual atoms, energy is released.
In the formation of chemical bonds, the electrons of the outer electron layer with the lowest binding energy with the nucleus, called valence. For example, in boron, these are electrons of the 2nd energy level - 2 electrons per 2 s- orbitals and 1 by 2 p-orbitals:
When a chemical bond is formed, each atom tends to obtain an electronic configuration of noble gas atoms, i.e. so that in its outer electron layer there are 8 electrons (2 for elements of the first period). This phenomenon is called the octet rule.
It is possible for atoms to achieve the electronic configuration of a noble gas if initially single atoms share some of their valence electrons with other atoms. In this case, common electron pairs are formed.
Depending on the degree of socialization of electrons, covalent, ionic and metallic bonds can be distinguished.
covalent bond
A covalent bond occurs most often between atoms of non-metal elements. If the atoms of non-metals forming a covalent bond belong to different chemical elements, such a bond is called a covalent polar bond. The reason for this name lies in the fact that atoms of different elements also have a different ability to attract a common electron pair to themselves. Obviously, this leads to a shift of the common electron pair towards one of the atoms, as a result of which a partial negative charge is formed on it. In turn, a partial positive charge is formed on the other atom. For example, in a hydrogen chloride molecule, the electron pair is shifted from the hydrogen atom to the chlorine atom:
Examples of substances with a covalent polar bond:
СCl 4 , H 2 S, CO 2 , NH 3 , SiO 2 etc.
A covalent non-polar bond is formed between non-metal atoms of the same chemical element. Since the atoms are identical, their ability to pull shared electrons is the same. In this regard, no displacement of the electron pair is observed:
The above mechanism for the formation of a covalent bond, when both atoms provide electrons for the formation of common electron pairs, is called exchange.
There is also a donor-acceptor mechanism.
When a covalent bond is formed by the donor-acceptor mechanism, a common electron pair is formed due to the filled orbital of one atom (with two electrons) and the empty orbital of another atom. An atom that provides an unshared electron pair is called a donor, and an atom with a free orbital is called an acceptor. The donors of electron pairs are atoms that have paired electrons, for example, N, O, P, S.
For example, according to the donor-acceptor mechanism, the fourth N-H covalent bond is formed in the ammonium cation NH 4 +:
In addition to polarity, covalent bonds are also characterized by energy. The bond energy is the minimum energy required to break a bond between atoms.
The binding energy decreases with increasing radii of the bound atoms. Since we know that atomic radii increase down the subgroups, we can, for example, conclude that the strength of the halogen-hydrogen bond increases in the series:
HI< HBr < HCl < HF
Also, the bond energy depends on its multiplicity - the greater the bond multiplicity, the greater its energy. The bond multiplicity is the number of common electron pairs between two atoms.
Ionic bond
An ionic bond can be considered as the limiting case of a covalent polar bond. If in a covalent-polar bond the common electron pair is partially shifted to one of the pair of atoms, then in the ionic one it is almost completely “given away” to one of the atoms. The atom that has donated an electron(s) acquires a positive charge and becomes cation, and the atom that took electrons from it acquires a negative charge and becomes anion.
Thus, an ionic bond is a bond formed due to the electrostatic attraction of cations to anions.
The formation of this type of bond is characteristic of the interaction of atoms of typical metals and typical nonmetals.
For example, potassium fluoride. A potassium cation is obtained as a result of the detachment of one electron from a neutral atom, and a fluorine ion is formed by attaching one electron to a fluorine atom:
Between the resulting ions, a force of electrostatic attraction arises, as a result of which an ionic compound is formed.
During the formation of a chemical bond, electrons from the sodium atom passed to the chlorine atom and oppositely charged ions were formed, which have a completed external energy level.
It has been established that electrons do not completely detach from the metal atom, but only shift towards the chlorine atom, as in a covalent bond.
Most binary compounds that contain metal atoms are ionic. For example, oxides, halides, sulfides, nitrides.
An ionic bond also occurs between simple cations and simple anions (F -, Cl -, S 2-), as well as between simple cations and complex anions (NO 3 -, SO 4 2-, PO 4 3-, OH -). Therefore, ionic compounds include salts and bases (Na 2 SO 4, Cu (NO 3) 2, (NH 4) 2 SO 4), Ca (OH) 2, NaOH).
metal connection
This type of bond is formed in metals.
The atoms of all metals have electrons on the outer electron layer that have a low binding energy with the atomic nucleus. For most metals, the loss of external electrons is energetically favorable.
In view of such a weak interaction with the nucleus, these electrons in metals are very mobile, and the following process continuously occurs in each metal crystal:
M 0 - ne - \u003d M n +, where M 0 is a neutral metal atom, and M n + cation of the same metal. The figure below shows an illustration of the ongoing processes.
That is, electrons “rush” along the metal crystal, detaching from one metal atom, forming a cation from it, joining another cation, forming a neutral atom. This phenomenon was called “electronic wind”, and the set of free electrons in the crystal of a non-metal atom was called “electron gas”. This type of interaction between metal atoms is called a metallic bond.
hydrogen bond
If a hydrogen atom in a substance is bonded to an element with a high electronegativity (nitrogen, oxygen, or fluorine), the substance is characterized by the phenomenon of hydrogen bonding.
Since a hydrogen atom is bonded to an electronegative atom, a partial positive charge is formed on the hydrogen atom, and a partial negative charge is formed on the electronegative atom. In this regard, electrostatic attraction becomes possible between the partially positively charged hydrogen atom of one molecule and the electronegative atom of another. For example, hydrogen bonding is observed for water molecules:
It is the hydrogen bond that explains the abnormally high melting point of water. In addition to water, strong hydrogen bonds are also formed in substances such as hydrogen fluoride, ammonia, oxygen-containing acids, phenols, alcohols, amines.
In which one of the atoms donated an electron and became a cation, and the other atom accepted an electron and became an anion.
The characteristic properties of a covalent bond - directionality, saturation, polarity, polarizability - determine the chemical and physical properties of compounds.
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 (non-polar - a diatomic molecule consists of identical atoms (H 2, Cl 2, N 2) and the electron clouds of each atom are distributed symmetrically with respect to these atoms; polar - a diatomic molecule consists of atoms of different chemical elements , and the general electron cloud shifts towards one of the atoms, thereby forming an asymmetry in the distribution of the electric charge in the molecule, generating a dipole moment of the molecule).
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 electron mobility. The polarity and polarizability of covalent bonds determine the reactivity of molecules with respect to polar reagents.
However, twice Nobel Prize winner L. Pauling pointed out that "in some molecules there are covalent bonds due to one or three electrons instead of a common pair." A single-electron chemical bond is realized in the molecular ion hydrogen H 2 + .
The molecular hydrogen ion H 2 + contains two protons and one electron. The single electron of the molecular system compensates for the electrostatic repulsion of two protons and keeps them at a distance of 1.06 Å (the length of the H 2 + chemical bond). The center of the electron density of the electron cloud of the molecular system is equidistant from both protons by the Bohr radius α 0 =0.53 A and is the center of symmetry of the molecular hydrogen ion H 2 + .
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A covalent bond is formed by a pair of electrons shared between two atoms, and these electrons must occupy two stable orbitals, one from each atom.
A + B → A: B
As a result of socialization, electrons form a filled energy level. A bond is formed if their total energy at this level is less than in the initial state (and the difference in energy will be nothing more than the bond energy).
According to the theory of molecular orbitals, the overlap of two atomic orbitals leads in the simplest case to the formation of two molecular orbitals (MOs): binding MO and antibonding (loosening) MO. Shared electrons are located on a lower energy binding MO.
Formation of a bond during the recombination of atoms
However, the mechanism of interatomic interaction remained unknown for a long time. Only in 1930, F. London introduced the concept of dispersion attraction - the interaction between instantaneous and induced (induced) dipoles. At present, the attractive forces due to the interaction between fluctuating electric dipoles of atoms and molecules are called "London forces".
The energy of such an interaction is directly proportional to the square of the electronic polarizability α and inversely proportional to the distance between two atoms or molecules to the sixth power.
Bond formation by the donor-acceptor mechanism
In addition to the homogeneous mechanism for the formation of a covalent bond described in the previous section, there is a heterogeneous mechanism - the interaction of oppositely charged ions - the proton H + and the negative hydrogen ion H -, called the hydride ion:
H + + H - → H 2
When the ions approach, the two-electron cloud (electron pair) of the hydride ion is attracted to the proton and eventually becomes common to both hydrogen nuclei, that is, it turns into a binding electron pair. The particle that supplies an electron pair is called a donor, and the particle that accepts this electron pair is called an acceptor. Such a mechanism for the formation of a covalent bond is called donor-acceptor.
H + + H 2 O → H 3 O +
A proton attacks the lone electron pair of a water molecule and forms a stable cation that exists in aqueous solutions of acids.
Similarly, a proton is attached to an ammonia molecule with the formation of a complex ammonium cation:
NH 3 + H + → NH 4 +
In this way (according to the donor-acceptor mechanism for the formation of a covalent bond), a large class of onium compounds is obtained, which includes ammonium, oxonium, phosphonium, sulfonium and other compounds.
A hydrogen molecule can act as an electron pair donor, which, upon contact with a proton, leads to the formation of a molecular hydrogen ion H 3 + :
H 2 + H + → H 3 +
The binding electron pair of the molecular hydrogen ion H 3 + belongs simultaneously to three protons.
Types of covalent bond
There are three types of covalent chemical bonds that differ in the mechanism of 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 that form a simple covalent bond are the same, then the true charges of the atoms in the molecule are also the same, since the atoms that form the bond equally own a shared electron pair. Such a connection is called non-polar covalent bond. Simple substances have such a connection, for example: 2, 2, 2. But not only non-metals of the same type can form a covalent non-polar bond. Non-metal elements whose electronegativity is of equal value can also form a covalent non-polar bond, for example, in the PH 3 molecule, the bond is covalent non-polar, since the EO of hydrogen is equal to the EO of phosphorus.
- 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 greater electronegativity attracts a pair of bond electrons to itself more strongly, and its true charge becomes negative. An atom with less electronegativity acquires, respectively, the same positive charge. If a compound is formed between two different non-metals, then such a compound is called polar covalent bond.
In the ethylene molecule C 2 H 4 there is a double bond CH 2 \u003d CH 2, its electronic formula: 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, weaker covalent bond is called π (\displaystyle \pi )-communication.
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 (\displaystyle \pi ) bonds between the same carbon atoms. Two π (\displaystyle \pi )-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 (\displaystyle \pi )-connections, but a single π (\displaystyle \pi ) dielectrics or semiconductors. Typical examples of atomic crystals (the atoms in which are interconnected by covalent (atomic) bonds) are
Covalent bond Formation mechanism according to Lewis.
A bond between atoms occurs when their atomic orbitals overlap to form molecular orbitals (MOs). There are two mechanisms for the formation of a covalent bond.
EXCHANGE MECHANISM - one-electron atomic orbitals participate in the formation of a bond, i.e. each of the atoms provides for the common use of one electron:
DONOR-ACCEPTOR MECHANISM - the formation of a bond occurs due to a pair of electrons of the donor atom and the vacant orbital of the acceptor atom: \\
The characteristics of a covalent bond do not depend on the mechanism of its formation.
Properties of a covalent bond: saturation, directionality, hybridization, multiplicity.
The features of a covalent bond are its directionality and saturation. Since atomic orbitals are spatially oriented, the overlap of electron clouds occurs in certain directions, which determines the direction of the covalent bond. Directivity is expressed quantitatively as bond angles between chemical bond directions in molecules and solids. Saturation of a covalent bond is caused by limiting the number of electrons in the outer shells that can participate in the formation of a covalent bond.
CS properties:
1. COP strength- these are the properties of the nature of the long bond (internuclear space) and the energy of the bond energy.
2. Polarity of the COP. In molecules containing atomic nuclei of the same element, one or more pairs of electrons equally belong to both atoms, each atomic nucleus attracts a pair of binding electrons with equal force. Such a connection is called non-polar covalent bond.
If a pair of electrons forming a chemical bond is shifted to one of the nuclei of atoms, then the bond is called polar covalent bond.
3. Saturation of the CS- this is the ability of an atom to participate only in a certain number of CSs, saturation characterizes the valency of the atom. Quantitative measures of valency yavl. the number of unpaired electrons in an atom in the ground and in the excited state.
4. Orientation of the COP. The strongest CSs are formed in the direction of the maximum overlap of atomic orbitals, i.e. The measure of direction is the bond angle.
5. Hybridization of CS - during hybridization, a shift of atomic orbitals occurs, i.e. there is an alignment in energy and in form. Exists sp, sp2, sp3 - hybridization. sp- the shape of the molecule is linear (angle 180 0), sp2- the shape of the molecule is flat triangular (angle 120 0) , sp 3 - tetrahedral shape (angle 109 0 28).
6. The multiplicity of the CS or the decolization of the connection The number of bonds formed between atoms is called multiplicity (order) connections. With an increase in the multiplicity (order) of the bond, the bond length and its energy change.
covalent bond(from the Latin "with" jointly and "vales" valid) is carried out by an electron pair belonging to both atoms. Formed between atoms of non-metals.
The electronegativity of non-metals is quite large, so that in the chemical interaction of two atoms of non-metals, the complete transfer of electrons from one to the other (as in the case) is impossible. In this case, electron pooling is necessary to perform.
As an example, let's discuss the interaction of hydrogen and chlorine atoms:
H 1s 1 - one electron
Cl 1s 2 2s 2 2 p6 3 s2 3 p5 - seven electrons in the outer level
Each of the two atoms lacks one electron in order to have a complete outer electron shell. And each of the atoms allocates “for common use” one electron. Thus, the octet rule is satisfied. The best way to represent this is with the Lewis formulas:
Formation of a covalent bond
The shared electrons now belong to both atoms. The hydrogen atom has two electrons (its own and the shared electron of the chlorine atom), and the chlorine atom has eight electrons (its own plus the shared electron of the hydrogen atom). These two shared electrons form a covalent bond between the hydrogen and chlorine atoms. The particle formed when two atoms bond is called molecule.
Non-polar covalent bond
A covalent bond can form between two the same atoms. For example:
This diagram explains why hydrogen and chlorine exist as diatomic molecules. Thanks to the pairing and socialization of two electrons, it is possible to fulfill the octet rule for both atoms.
In addition to single bonds, a double or triple covalent bond can be formed, as, for example, in oxygen O 2 or nitrogen N 2 molecules. Nitrogen atoms each have five valence electrons, so three more electrons are required to complete the shell. This is achieved by sharing three pairs of electrons, as shown below:
Covalent compounds are usually gases, liquids, or relatively low-melting solids. One of the rare exceptions is diamond, which melts above 3,500°C. This is due to the structure of diamond, which is a continuous lattice of covalently bonded carbon atoms, and not a collection of individual molecules. In fact, any diamond crystal, regardless of its size, is one huge molecule.
A covalent bond occurs when the electrons of two nonmetal atoms join together. The resulting structure is called a molecule.
Polar covalent bond
In most cases, two covalently bonded atoms have different electronegativity and shared electrons do not belong equally to two atoms. Most of the time they are closer to one atom than to another. In a molecule of hydrogen chloride, for example, the electrons that form a covalent bond are located closer to the chlorine atom, since its electronegativity is higher than that of hydrogen. However, the difference in the ability to attract electrons is not so great that there is a complete transfer of an electron from a hydrogen atom to a chlorine atom. Therefore, the bond between hydrogen and chlorine atoms can be viewed as a cross between an ionic bond (complete electron transfer) and a non-polar covalent bond (symmetrical arrangement of a pair of electrons between two atoms). The partial charge on atoms is denoted by the Greek letter δ. Such a connection is called polar covalent bond, and the hydrogen chloride molecule is said to be polar, that is, it has a positively charged end (hydrogen atom) and a negatively charged end (chlorine atom).
The table below lists the main types of bonds and examples of substances:
Exchange and donor-acceptor mechanism of covalent bond formation
1) Exchange mechanism. Each atom contributes one unpaired electron to a common electron pair.
2) Donor-acceptor mechanism. One atom (donor) provides an electron pair, and another atom (acceptor) provides an empty orbital for this pair.
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