Covalent bond

The characteristic properties of the covalent bond — directivity, saturation, polarity, polarizability — determine the chemical and physical properties of the compounds.

• The orientation 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 valence.

• 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 bond polarity is due to the uneven distribution of electron density due to differences in the electronegativity of atoms.

According to this criterion, 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 relative to these atoms; polar - a diatomic molecule consists of atoms of different chemical elements , and the general electron cloud shifts toward one of the atoms, thereby forming an asymmetry in the distribution of electric charge in the molecule, generating the dipole moment of the molecule).

• The polarizability of the bond is expressed in the displacement of the bond electrons under the influence of an external electric field, including another reacting particle. The polarizability is determined by the mobility of electrons . The polarity and polarizability of covalent bonds determines the reactivity of molecules with respect to polar reagents.

Electrons are more mobile, the farther they are from the nuclei.

However, twice the Nobel laureate L. Pauling pointed out that “in some molecules there are covalent bonds due to one or three electrons instead of a common pair” ^[4] . A one-electron chemical bond is realized in the molecular hydrogen ion H ₂ ⁺ .

The molecular hydrogen ion H ₂ ⁺ contains two protons and one electron. The only electron in the molecular system compensates for the electrostatic repulsion of two protons and holds them at a distance of 1.06 Å ( chemical bond length H ₂ ⁺ ). The center of 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 ⁺ .

The term "covalent bond" was first coined by Nobel laureate Irving Langmuir in 1919 ^{[5] [4]} . This term refers to a chemical bond due to the joint possession of electrons , as opposed to a metal bond in which the electrons were free, or from an ionic bond in which one of the atoms gave away an electron and became a cation , and the other atom took 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 covalent bonds from the point of view of quantum mechanics .

A covalent bond is formed by a pair of electrons divided between two atoms, and these electrons must occupy two stable orbitals, one from each atom ^[6] .

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 binding energy ).

According to the theory of molecular orbitals, the overlap of two atomic orbitals in the simplest case leads to the formation of two molecular orbitals (MO): a binding MO and an anti- binding (loosening) MO . Socialized electrons are located at a lower energy binding MO.

Atom recombination bonding

Atoms and free radicals are prone to recombination - the formation of a covalent bond by the socialization of two unpaired electrons belonging to different particles.

H + H → H ₂ ;

CH ₃ + CH ₃ → CH ₃ - CH ₃ .

The formation of bonds during recombination is accompanied by the release of energy. So, in the interaction of hydrogen atoms , energy is released in the amount of 436 kJ / mol. This effect is used in technology for atomic hydrogen welding. A stream of hydrogen is passed through an electric arc, where a stream of hydrogen atoms is generated. Atoms are then reunited on a metal surface placed a short distance from the arc. Metal can be heated in this way above 3500 ° C. The great advantage of the “atomic hydrogen flame” is the uniformity of heating, which allows welding of very thin metal parts ^[7] .

However, the mechanism of interatomic interaction has long remained unknown. Only in 1930, F. London introduced the concept of dispersion attraction - the interaction between instantaneous and induced (induced) dipoles. Currently, the attractive forces due to the interaction between fluctuating electric dipoles of atoms and molecules are called " London forces ".

The energy of this interaction is directly proportional to the square of the electron polarizability α and is inversely proportional to the distance between two atoms or molecules to the sixth degree ^[8] .

Donor-acceptor bond formation

In addition to the homogeneous mechanism for the formation of covalent bonds described in the previous section, there is a heterogeneous mechanism - the interaction of oppositely charged ions - the H ⁺ proton and the negative hydrogen ion H ^- , called the hydride ion :

H ⁺ + H ^- → H ₂

When ions approach each other, the two-electron cloud (electron pair) of the hydride ion is attracted to the proton and ultimately becomes common to both hydrogen nuclei, that is, it turns into a binding electron pair. A particle supplying an electron pair is called a donor, and a particle receiving this electron pair is called an acceptor. This mechanism of covalent bond formation is called donor – acceptor ^[9] .

The distribution of electron density between nuclei in a hydrogen molecule is one and the same, regardless of the formation mechanism; therefore, it is incorrect to call a chemical bond obtained by the donor – acceptor mechanism a donor – acceptor bond.

In addition to the hydride ion, the donors of the electron pair are compounds of elements of the main subgroups of groups V – VII of the periodic system of elements to the lowest degree of element oxidation. So, even Johannes Bronsted found that the proton does not exist in solution in its free form, in water it forms an oxonium cation:

H ⁺ + H ₂ O → H ₃ O ⁺

A proton attacks an unshared electron pair of a water molecule and forms a stable cation that exists in aqueous solutions of acids ^[10] .

Similarly, the proton joins the ammonia molecule with the formation of a complex ammonium cation:

NH ₃ + H ⁺ → NH ₄ ⁺

In this way (by the donor – acceptor mechanism of covalent bond formation), a large class of onium compounds is obtained, which includes ammonium , oxonium, phosphonium, sulfonium, and other compounds ^[11] .

A hydrogen molecule can act as a donor of an electron pair, 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.

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 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 possess a socialized electron pair. Such a bond is called a nonpolar covalent bond . This connection has many simple substances , for example: O ₂ , N ₂ , Cl ₂ . But not only non-metals of the same type can form a covalent non-polar bond. Non-metallic elements can also form a covalent non-polar bond, the electronegativity of which is equally important, for example, in a 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, accordingly, the same positive charge. If a compound is formed between two different non-metals , then such a compound is called a covalent polar bond .

2. Donor-acceptor communication . To form this type of covalent bond, both electrons are provided by one of the atoms - the donor . The second of the atoms involved in the formation of a bond is called an acceptor . In the resulting molecule, the formal charge of the donor increases by one, and the formal charge of the acceptor decreases by one.

3. Semipolar communication . It can be considered as a polar donor-acceptor bond. This type of covalent bond is formed between an atom having an unshared pair of electrons ( nitrogen , phosphorus , sulfur , halogens , etc.) and an atom with two unpaired electrons ( oxygen , sulfur ). The formation of a semipolar connection proceeds in two stages:

1. Transfer of one electron from an atom with an unshared pair of electrons to an atom with two unpaired electrons. As a result, an atom with an unshared electron pair transforms into a radical cation (a positively charged particle with an unpaired electron), and an atom with two unpaired electrons turns into a radical anion (a negatively charged particle with an unpaired electron).

2. Socialization of unpaired electrons (as in the case of a simple covalent bond).

In the formation of a semipolar bond, an atom with an unshared pair of electrons increases its formal charge by one, and an atom with two unpaired electrons lowers its formal charge by one.

Sigma (σ) - , pi (π) bonds - an approximate description of the types of covalent bonds in the molecules of various 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. In education${\displaystyle \mathrm{<span\; class="mwe-math-element"><span\; class="mwe-math-mathml-inline\; mwe-math-mathml-a11y"\; style="display:\; none;">\pi </span></span>}}$ {\ displaystyle \ pi}   -connection is carried out the so-called lateral overlapping of electronic clouds, 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 = 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 the σ-bond; the second, less strong covalent bond is called${\displaystyle \mathrm{<span\; class="mwe-math-element"><span\; class="mwe-math-mathml-inline\; mwe-math-mathml-a11y"\; style="display:\; none;">\pi </span></span>}}$ {\ displaystyle \ pi}   -connection.

In a linear acetylene molecule

H — C≡C — H (H: C ::: C: H)

there are σ bonds between carbon and hydrogen atoms, one σ bond between two carbon atoms and two${\displaystyle \mathrm{<span\; class="mwe-math-element"><span\; class="mwe-math-mathml-inline\; mwe-math-mathml-a11y"\; style="display:\; none;">\pi </span></span>}}$ {\ displaystyle \ pi}   - bonds between the same carbon atoms. Two${\displaystyle \mathrm{<span\; class="mwe-math-element"><span\; class="mwe-math-mathml-inline\; mwe-math-mathml-a11y"\; style="display:\; none;">\pi </span></span>}}$ {\ displaystyle \ pi}   -bonds are located above the scope of the σ-bond in two mutually perpendicular planes.

All six carbon atoms of the cyclic C ₆ H ₆ benzene molecule lie in the same plane. Between the carbon atoms in the plane of the ring are σ-bonds; each carbon atom with hydrogen atoms has the same bonds. Carbon atoms spend three electrons each to make these bonds. Clouds of the fourth valence electrons of carbon atoms, in the form of eights, are located perpendicular to the plane of the benzene molecule. Each such cloud overlaps identically with the electron clouds of neighboring carbon atoms. In the benzene molecule, not three separate${\displaystyle \mathrm{<span\; class="mwe-math-element"><span\; class="mwe-math-mathml-inline\; mwe-math-mathml-a11y"\; style="display:\; none;">\pi </span></span>}}$ {\ displaystyle \ pi}   -connection, and a single${\displaystyle \mathrm{<span\; class="mwe-math-element"><span\; class="mwe-math-mathml-inline\; mwe-math-mathml-a11y"\; style="display:\; none;">\pi </span></span>}}$ {\ displaystyle \ pi}   -electronic system of six electrons, common to all carbon atoms. The bonds between the carbon atoms in the benzene molecule are exactly the same.

A simple covalent bond connects the atoms in the molecules of simple gases (H ₂ , Cl ₂ , etc.) and compounds (H ₂ O, NH ₃ , CH ₄ , CO ₂ , HCl, etc.). Compounds with a donor – acceptor bond — ammonium NH ₄ ⁺ , tetrafluoroborate anion BF ₄ ^— and others. Compounds with a semipolar bond — nitrous oxide N ₂ O, O ^— -PCl ₃ ⁺ .

Covalently bonded crystals are dielectrics or semiconductors . Typical examples of atomic crystals (atoms in which are interconnected by covalent (atomic) bonds) can serve as diamond , germanium and silicon .

• Donor-acceptor bond
• Chemical bond polarization
• Ionic bond
• Metal bond
• Recombination

• «Химический энциклопедический словарь», М., «Советская энциклопедия», 1983, с.264.