Types of hybridization. Hybridization of electron orbitals and molecular geometry

Hybridization atomic orbitals- a process that allows us to understand how atoms modify their orbitals when forming compounds. So, what is hybridization, and what types of it exist?

General characteristics of hybridization of atomic orbitals

Atomic orbital hybridization is a process in which different orbitals of a central atom are mixed, resulting in the formation of orbitals with identical characteristics.

Hybridization occurs during the formation of a covalent bond.

The hybrid orbital has the odds of an infinity sign or an asymmetrical inverted figure of eight, extended away from the atomic nucleus. This form causes a stronger overlap of hybrid orbitals with the orbitals (pure or hybrid) of other atoms than in the case of pure atomic orbitals and leads to the formation of stronger covalent bonds.

Rice. 1. Hybrid orbital appearance.

The idea of hybridization of atomic orbitals was first put forward by the American scientist L. Pauling. He believed that an atom entering into a chemical bond has different atomic orbitals (s-, p-, d-, f-orbitals), and as a result, hybridization of these orbitals occurs. The essence of the process is that atomic orbitals equivalent to each other are formed from different orbitals.

Types of atomic orbital hybridization

There are several types of hybridization:

. This type of hybridization occurs when one s orbital and one p orbital mix. As a result, two full-fledged sp orbitals are formed. These orbitals are located towards the atomic nucleus in such a way that the angle between them is 180 degrees.

Rice. 2. sp-hybridization.

sp2 hybridization. This type of hybridization occurs when one s orbital and two p orbitals mix. As a result, three hybrid orbitals are formed, which are located in the same plane at an angle of 120 degrees to each other.
. This type of hybridization occurs when one s orbital and three p orbitals mix. As a result, four full-fledged sp3 orbitals are formed. These orbitals are directed towards the top of the tetrahedron and are located at an angle of 109.28 degrees to each other.

sp3 hybridization is characteristic of many elements, for example, the carbon atom and other substances of group IV (CH 4, SiH 4, SiF 4, GeH 4, etc.)

Rice. 3. sp3 hybridization.

More complex types of hybridization involving d-orbitals of atoms are also possible.

What have we learned?

Hybridization is complex chemical process, when different orbitals of an atom form identical (equivalent) hybrid orbitals. The theory of hybridization was first voiced by the American L. Pauling. There are three main types of hybridization: sp-hybridization, sp2-hybridization, sp3-hybridization. There are also more complex types of hybridization that involve d orbitals.

One of the tasks of chemistry is the study of the structure of matter, including elucidation of the mechanism of formation of various compounds from simple substances formed by atoms of one chemical element. Features of the interaction of atoms, more precisely, their differently charged components - electronic shells and nuclei - are described as different types of chemical bonds. Thus, substances are formed through covalent bonds, to describe which in 1931 the American chemist L. Pauling proposed a model of hybridization of atomic orbitals.

Concept of covalent bond

In cases where the interaction process results in the formation of a pair of valence electron clouds common to two atoms, we speak of a covalent bond. As a result of its occurrence, smallest particle simple or complex substance - molecule.

One of the features of a covalent bond is its directionality - a consequence complex shape electron orbitals p, d and f, which, without having spherical symmetry, have a certain spatial orientation. Another one important feature of this type of chemical bond - saturation, due to the limited number of external - valence - clouds in the atom. That is why the existence of a molecule, for example, H 2 O, is possible, but H 5 O is not.

Types of covalent bond

The formation of shared electron pairs can occur different ways. In the mechanism of covalent bond formation important role plays a role in the nature of cloud overlap and the spatial symmetry of the resulting cloud. According to this criterion, L. Pauling proposed to distinguish the following types:

The sigma bond (σ) has the greatest degree of overlap along the axis passing through atomic nuclei. Here the cloud density will be maximum.
The pi bond (π) is formed by lateral overlap, and the electron cloud, accordingly, has the highest density outside the axis connecting the nuclei.

These spatial characteristics are of great importance insofar as they correlate with the energy parameters of the covalent bond.

Features of polyatomic molecules

The concept of hybridization was introduced by Pauling to explain one of the features of covalent bonds in polyatomic molecules. It is known that the bonds formed by the central atom in such molecules turn out to be identical in spatial and energy characteristics. This occurs regardless of which orbitals (s, p or d) are involved in the formation of a common electron pair.

Very convenient and a clear example The carbon atom is used to illustrate this phenomenon. When entering into a chemical bond, an atom in an excited state has 4 valence orbitals: 2s, 2p x, 2p y and 2p z. The last three differ from the 2s orbital in energy and shape. Nevertheless, in a molecule of, for example, CH4 methane, all four bonds are completely equivalent and have bond angles of 109.5° (while the p-orbitals are located at angles of 90°). In other carbon compounds, bond angles of 120° and 180° occur; in molecules containing nitrogen (ammonia NH 3) and oxygen (water H 2 O) these angles are 107.5° and 104.5°. The appearance of such bond angles also required explanation.

The essence of the phenomenon

The idea of hybridization is the formation of averaged orbitals by overlapping electron clouds different types with close energy values - s, p, sometimes d. The number of resulting - hybrid - orbitals corresponds to the number of overlapping clouds. Since an orbital is the determining probability of finding an electron at a particular point in an atom, a hybrid orbital is a superposition of wave functions that occurs as a result of electronic transitions when the atom is excited. It leads to the emergence of equivalent wave functions that differ only in direction.

Hybrid orbitals are equivalent in energy and have the same shape in the form of a three-dimensional figure eight, which has a strong asymmetry relative to the nucleus. Less energy is spent on hybridization than is released during the formation of a strong covalent bond with hybrid orbitals, therefore this process is energetically favorable, that is, the most probable.

orbital hybridization and molecular geometry

Possible various options overlap (mixing) of external electron clouds in an atom. The most common types of orbital superposition are:

Sp 3 -hybridization. This option is realized by superposing one s- and three p-orbitals. The result is four hybrid orbitals, the axes of which are directed for any pair at angles of 109.5°, corresponding to the minimum mutual repulsion of electrons. When these orbitals enter into σ bonds with other atoms, a molecule of a tetrahedral configuration is formed, for example, methane, ethane C 2 H 6 (a combination of two tetrahedra), ammonia, water. In an ammonia molecule, one, and in a water molecule, two of the vertices of the tetrahedron are occupied by lone electron pairs, which leads to a decrease in the bond angle.
Sp 2 hybridization occurs when one s and two p orbitals are combined. In this case, the three hybrid orbitals are located at angles of 120° in the same plane. Similar triangular shape have, for example, molecules of boron trichloride BCl 3, which is used in various technologies. Another example, the ethylene molecule, is formed due to an additional π bond between carbon atoms, in which one p orbital is non-hybrid and oriented perpendicular to the plane formed by two triangles.
Sp hybridization occurs when one s and one p orbital mix. The two hybrid clouds are located at an angle of 180°, and the molecule has a linear configuration. Examples are molecules of beryllium chloride BeCl 2 or acetylene C 2 H 2 (in the latter, two non-hybrid carbon p-orbitals form additional π bonds).

There are also more complex options for the hybridization of atomic orbitals: sp 3 d, sp 3 d 2 and others.

The role of the hybridization model

Pauling's concept provides a good qualitative description of the structure of molecules. It is convenient and visual, and successfully explains some of the features of covalent compounds, such as the size of bond angles or the alignment of the length of a chemical bond. However, the quantitative side of the model cannot be considered satisfactory, since it does not allow making many important predictions regarding physical effects associated with the structural features of molecules, for example, molecular photoelectron spectra. The author of the concept of hybridization himself already noted its shortcomings in the early 1950s.

Nevertheless, the model of hybridization of atomic orbitals played a major role in the development of modern ideas about the structure of matter. Based on it, more adequate concepts were developed, for example, the theory of repulsion of electron pairs. Therefore, of course, the hybridization model was an important stage in the development of theoretical chemistry, and when describing some aspects of the electronic structure of molecules, it is quite applicable at the present time.

The valence bond method makes it possible to clearly explain the spatial characteristics of many molecules. However, the usual idea of the shapes of orbitals is not enough to answer the question why, if the central atom has different - s, p, d– valence orbitals, the bonds formed by it in molecules with identical substituents turn out to be equivalent in their energy and spatial characteristics. In the twenties of the 19th century, Linus Pauling proposed the concept of hybridization of electron orbitals. Hybridization is an abstract model of the alignment of atomic orbitals in shape and energy.

Examples of hybrid orbital shapes are presented in Table 5.

Table 5. Hybrid sp, sp 2 , sp 3 orbitals

The concept of hybridization is convenient to use when explaining the geometric shape of molecules and the size of bond angles (examples of tasks 2–5).

Algorithm for determining the geometry of molecules using the BC method:

A. Determine the central atom and the number of σ-bonds with the terminal atoms.

b. Draw up the electronic configurations of all atoms that make up the molecule and graphic images of the outer electronic levels.

V. According to the principles of the BC method, the formation of each bond requires a pair of electrons, in the general case, one from each atom. If there are not enough unpaired electrons for the central atom, one should assume the excitation of the atom with the transition of one of the pair of electrons to a higher energy level.

d. Assume the need and type of hybridization, taking into account all bonds and, for elements of the first period, unpaired electrons.

e. Based on the above conclusions, draw the electronic orbitals (hybrid or not) of all atoms in the molecule and their overlap. Draw a conclusion about the geometry of the molecule and the approximate value of bond angles.

f. Determine the degree of bond polarity based on the electronegativity values of the atoms (Table 6) Determine the presence of a dipole moment based on the location of the centers of gravity of positive and negative charges and/or symmetry of the molecule.

Table 6. Electronegativity values of some elements according to Pauling

Examples of tasks

Exercise 1. Describe the chemical bond in the CO molecule using the BC method.

Solution (Fig. 25)

A. Draw up the electronic configurations of all the atoms that make up the molecule.

b. To form a bond, it is necessary to create socialized electron pairs

Figure 25. Scheme of bond formation in a CO molecule (without orbital hybridization)

Conclusion: In the CO molecule there is a triple bond C≡O

For the CO molecule, we can assume the presence sp-hybridization of the orbitals of both atoms (Fig. 26). Paired electrons not involved in bond formation are located on sp-hybrid orbital.

Figure 26. Scheme of bond formation in a CO molecule (taking into account the hybridization of orbitals)

Task 2. Based on the BC method, assume the spatial structure of the BeH 2 molecule and determine whether the molecule is a dipole.

The solution to the problem is presented in Table 7.

Table 7. Determination of the geometry of the BeH 2 molecule

	Electronic configuration		Notes
A.			The central atom is beryllium. It needs to form two ϭ-bonds with hydrogen atoms
b.	H: 1 s 1 Be: 2 s 2		The hydrogen atom has an unpaired electron, the beryllium atom has all its electrons paired, it must be transferred to an excited state
V.	H: 1 s 1 Be*: 2 s 1 2p 1		If one hydrogen atom bonded with beryllium due to 2 s-electron of beryllium, and the other - due to 2 p-electron of beryllium, then the molecule would not have symmetry, which is not energetically justified, and the Be–H bonds would not be equivalent.
G.	H: 1 s 1 Be*: 2( sp) 2		It should be assumed that there is sp-hybridization
d.			Two sp-hybrid orbitals are located at an angle of 180°, the BeH 2 molecule is linear
e.		Electronegativity χ H = 2.1, χ Be = 1.5, therefore the bond is covalent polar, the electron density is shifted to the hydrogen atom, a small negative charge δ– appears on it.

On the beryllium atom δ+. Since the centers of gravity of the positive and negative charge coincide (it is symmetrical), the molecule is not a dipole. sp Similar reasoning will help describe the geometry of molecules with 2 - and sp

3-hybrid orbitals (Table 8).

Table 8. Geometry of BF 3 and CH 4 molecules Task 3.

Based on the BC method, assume the spatial structure of the H 2 O molecule and determine whether the molecule is a dipole. Two solutions are possible; they are presented in tables 9 and 10.

	Electronic configuration	Table 9. Determination of the geometry of the H 2 O molecule (without orbital hybridization)	Notes
A.
b.	H: 1 s Graphic representation of outer level orbitals s 2 2p 4
V.			1 O: 2
G.			Hybridization can be neglected
d.
e.

Thus, a water molecule should have a bond angle of about 90°. However, the angle between the bonds is approximately 104°.

This can be explained

1) repulsion of hydrogen atoms located close to each other.

2) Hybridization of orbitals (Table 10).

Table 10. Determination of the geometry of the H 2 O molecule (taking into account the hybridization of orbitals)

	Electronic configuration	Table 9. Determination of the geometry of the H 2 O molecule (without orbital hybridization)	Notes
A.			The central atom is oxygen. It needs to form two ϭ bonds with hydrogen atoms.
b.	H: 1 s Graphic representation of outer level orbitals s 2 2p 4		A hydrogen atom has an unpaired electron, and an oxygen atom has two unpaired electrons.
V.			A hydrogen atom has an unpaired electron, and an oxygen atom has two unpaired electrons.
G.			An angle of 104° suggests the presence sp 3-hybridization.
d.			Two sp The 3-hybrid orbitals are located at an angle of approximately 109°, the H 2 O molecule is close in shape to a tetrahedron, the decrease in the bond angle is explained by the influence of the electron non-bonding pair.
e.		Electronegativity χ Н = 2.1, χ О = 3.5, therefore the bond is covalent polar, the electron density is shifted to the oxygen atom, a small negative charge 2δ– appears on it. On the hydrogen atom δ+. Since the centers of gravity of the positive and negative charges do not coincide (it is not symmetrical), the molecule is a dipole.

Similar reasoning allows one to explain the bond angles in the ammonia molecule NH 3 . Hybridization involving lone electron pairs is usually assumed only for the orbitals of atoms of period II elements. Bond angles in molecules H 2 S = 92°, H 2 Se = 91°, H 2 Te = 89°. The same is observed in the series NH 3, РH 3, AsH 3. When describing the geometry of these molecules, traditionally, either they do not resort to the concept of hybridization, or they explain the decrease in the tetrahedral angle by the increasing influence of the lone pair.

Hybridization of atomic orbitals and molecular geometry

Important characteristic a molecule consisting of more than two atoms is its geometric configuration. It is determined by the mutual arrangement of atomic orbitals involved in the formation of chemical bonds.

Overlapping of electron clouds is possible only with a certain relative orientation of the electron clouds; in this case, the overlap region is located in a certain direction with respect to the interacting atoms.

Table 1 Hybridization of orbitals and spatial configuration of molecules

An excited beryllium atom has a configuration of 2s 1 2p 1, an excited boron atom has a configuration of 2s 1 2p 2, and an excited carbon atom has a configuration of 2s 1 2p 3. Therefore, we can assume that not the same, but different atomic orbitals can participate in the formation of chemical bonds. For example, in compounds such as BeCl 2, BeCl 3, CCl 4 there should be bonds of unequal strength and direction, and σ-bonds from p-orbitals should be stronger than bonds from s-orbitals, because for p-orbitals there are more favorable conditions for overlapping. However, experience shows that in molecules containing central atoms with different valence orbitals (s, p, d), all bonds are equivalent. An explanation for this was given by Slater and Pauling. They concluded that different orbitals, not very different in energy, form a corresponding number of hybrid orbitals. Hybrid (mixed) orbitals are formed from different atomic orbitals. The number of hybrid orbitals is equal to the number of atomic orbitals involved in hybridization. Hybrid orbitals are identical in electron cloud shape and energy. Compared to atomic orbitals, they are more elongated in the direction of formation of chemical bonds and therefore provide better overlap of electron clouds.

Hybridization of atomic orbitals requires energy, so hybrid orbitals in an isolated atom are unstable and tend to turn into pure AOs. When chemical bonds are formed, the hybrid orbitals are stabilized. Due to the stronger bonds formed by the hybrid orbitals, more energy is released from the system and therefore the system becomes more stable.

sp-hybridization occurs, for example, during the formation of Be, Zn, Co and Hg (II) halides. In the valence state, all metal halides contain s and p-unpaired electrons at the appropriate energy level. When a molecule is formed, one s and one p orbital form two hybrid sp orbitals at an angle of 180 degrees.

Fig.3 sp hybrid orbitals

Experimental data show that Be, Zn, Cd and Hg(II) halides are all linear and both bonds are of the same length.

sp 2 hybridization

As a result of the hybridization of one s-orbital and two p-orbitals, three hybrid sp 2 orbitals are formed, located in the same plane at an angle of 120 o to each other. This is, for example, the configuration of the BF 3 molecule:

Fig.4 sp 2 hybridization

sp 3 hybridization

sp 3 hybridization is characteristic of carbon compounds. As a result of the hybridization of one s orbital and three

p-orbitals, four hybrid sp 3 orbitals are formed, directed towards the vertices of the tetrahedron with an angle between the orbitals of 109.5 o. Hybridization is manifested in the complete equivalence of the bonds of a carbon atom with other atoms in compounds, for example, in CH 4, CCl 4, C(CH 3) 4, etc.

Fig.5 sp 3 hybridization

If all hybrid orbitals are connected to the same atoms, then the bonds are no different from each other. In other cases, slight deviations from standard bond angles occur. For example, in the water molecule H 2 O, oxygen - sp 3 -hybrid, is located in the center of an irregular tetrahedron, at the vertices of which two hydrogen atoms and two lone pairs of electrons “look” (Fig. 2). The shape of the molecule is angular when viewed from the centers of the atoms. The HOH bond angle is 105°, which is quite close to the theoretical value of 109°.

Fig.6 sp 3 - hybridization of oxygen and nitrogen atoms in molecules a) H 2 O and b) NCl 3.

If there were no hybridization (“alignment” of O-H bonds), the HOH bond angle would be 90° because the hydrogen atoms would be attached to two mutually perpendicular p orbitals. In this case, our world would probably look completely different.

The hybridization theory explains the geometry of the ammonia molecule. As a result of the hybridization of the 2s and three 2p orbitals of nitrogen, four sp 3 hybrid orbitals are formed. The configuration of the molecule is a distorted tetrahedron, in which three hybrid orbitals participate in the formation of a chemical bond, but the fourth with a pair of electrons does not. Angles between N-H bonds not equal to 90° as in a pyramid, but also not equal to 109.5°, corresponding to a tetrahedron.

Fig.7 sp 3 - hybridization in an ammonia molecule

When ammonia interacts with a hydrogen ion, as a result of donor-acceptor interaction, an ammonium ion is formed, the configuration of which is a tetrahedron.

Hybridization also explains the difference in angle between O-H connections in the corner water molecule. As a result of the hybridization of the 2s and three 2p orbitals of oxygen, four sp 3 hybrid orbitals are formed, of which only two are involved in the formation of a chemical bond, which leads to a distortion of the angle corresponding to the tetrahedron.

Fig.8 sp 3 hybridization in a water molecule

Hybridization can involve not only s- and p-orbitals, but also d- and f-orbitals.

With sp 3 d 2 hybridization, 6 equivalent clouds are formed. It is observed in such compounds as 4-, 4-. In this case, the molecule has the configuration of an octahedron:

Rice. 9 d 2 sp 3 -hybridization in ion 4-

Ideas about hybridization make it possible to understand such structural features of molecules that cannot be explained in any other way.

Hybridization of atomic orbitals (AO) leads to a displacement of the electron cloud in the direction of forming bonds with other atoms. As a result, the overlap areas of hybrid orbitals turn out to be larger than for pure orbitals and the bond strength increases.

Covalent bonds are the most common in the world of organic substances; they are characterized by saturation, polarizability and directionality in space.

The saturation of a covalent bond lies in the fact that the number of common electron pairs that a particular atom can form is limited. Due to this, covalent compounds have a strictly defined composition. Therefore, for example, there are molecules H 2, N 2, CH 4, but there are no molecules H 3, N 4, CH 5.

The polarizability of a covalent bond is the ability of molecules (and individual bonds in them) to change their polarity under the influence of external electric field- polarize.

As a result of polarization, nonpolar molecules can become polar, and polar ones can turn into even more polar ones, up to the complete breaking of individual bonds with the formation of ions:

The directionality of the covalent bond is due to the fact that the p-, d- and f-clouds are oriented in a certain way in space. The direction of a covalent bond affects the shape of the molecules of substances, their sizes, interatomic distances, bond angle, i.e., the geometry of the molecules.

A more complete picture of the shape of molecules of organic and inorganic substances can be formed on the basis of the hypothesis of hybridization of atomic orbitals. It was proposed by L. Pauling (USA) to explain what was established using physical methods studies of substances, the fact of the equivalence of all chemical bonds and their symmetrical arrangement relative to the center of the molecules CH 4, BF 3, BeCl 2. In each case, the formation of σ bonds from the central atom (C, B, Be) should involve electrons located in different states(s and p), so they could not be equivalent. The theory turned out to be unable to explain the facts; a contradiction arose, which was resolved with the help of a new hypothesis. This is one of the examples showing the path of development of human knowledge of the surrounding world, the possibility of ever deeper penetration into the essence of phenomena.

You became familiar with the hypothesis of hybridization of atomic orbitals in the course organic chemistry using the carbon atom as an example. Let us remind you of this again.

When a methane molecule CH 4 is formed, the carbon atom goes from the ground state to the excited one:

The outer electronic layer of the excited carbon atom contains one s-electron and three unpaired p-electrons, which form four σ-bonds with four s-electrons of hydrogen atoms. In this case, it should be expected that three C--H bonds formed due to the pairing of three p-electrons of a carbon atom with three s-electrons of three hydrogen atoms (s-p σ bond) should differ from the fourth (s-s) bond in strength , length, direction. A study of the electron density in methane molecules shows that all bonds in its molecule are equivalent and directed towards the vertices of the tetrahedron (Fig. 10). According to the hypothesis of hybridization of atomic orbitals, the four covalent bonds of the methane molecule are formed not with the participation of “pure” s- and p-clouds of the carbon atom, but with the participation of so-called hybrid, i.e., averaged, equivalent electron clouds.

Rice. 10. Ball-and-stick model of the methane molecule

According to this model, the number of hybrid atomic orbitals is equal to the number of the original “pure” orbitals. The corresponding hybrid clouds have a more favorable geometric shape than s- and p-clouds; their electron density is distributed differently, which ensures a more complete overlap with s-clouds of hydrogen atoms than would be the case for “pure” s- and p-clouds.

In the methane molecule and in other alkanes, as well as in all molecules of organic compounds, at the site of a single bond, carbon atoms are in a state of sp 3 hybridization, i.e., at the carbon atom, one s- and three p-atomic clouds underwent hybridization and four were formed identical hybrid sp 3 -atomic orbitals of the cloud.

As a result of the overlap of the corresponding four hybrid sp 3 clouds of the carbon atom with the s clouds of four hydrogen atoms, a tetrahedral methane molecule with four identical σ bonds located at an angle of 109°28" is formed (Fig. 11).

Rice. eleven.
Schemes of sp 3 hybridization of valence electron clouds (a) and the formation of bonds in a methane molecule (b)

This type of atomic hybridization and, consequently, the tetrahedral structure will also characterize the molecules of compounds of the carbon analogue - silicon: SiH 4, SiCl 4.

During the formation of water and ammonia molecules, sp 3 hybridization of the valence atomic orbitals of oxygen and nitrogen atoms also occurs. However, if the carbon atom has all four hybrid sp 3 clouds occupied by common electron pairs, then the nitrogen atom has one sp 3 cloud occupied by a lone electron pair, and the oxygen atom already has two sp 3 clouds occupied by them (Fig. 12).

Rice. 12.
Shapes of ammonia, water and hydrogen fluoride molecules

The presence of lone electron pairs leads to a decrease in bond angles (Table 8) compared to tetrahedral ones (109°28").

Table 8
Relationship between the number of lone electron pairs and the bond angle in molecules

sp 3 -Hybridization is observed not only for atoms in complex substances, but also for atoms in simple substances. For example, in atoms of such an allotropic modification of carbon as diamond.

In the molecules of some boron compounds, sp 2 hybridization of the valence atomic orbitals of the boron atom takes place.

For a boron atom in an excited state, one s- and two p-orbitals participate in hybridization, resulting in the formation of three sp 2 hybrid orbitals; the axes of the corresponding hybrid clouds are located in the plane at an angle of 120° to each other (Fig. 13).

Rice. 13.
Schemes of 8р 2 -hybridization and location of sp 2 -clouds in space

Therefore, the molecules of such compounds, for example BF3, have the shape of a flat triangle (Fig. 14).

Rice. 14.
Structure of the BF3 molecule

IN organic compounds, as you know, sp 2 hybridization is characteristic of carbon atoms in alkene molecules at the site of the double bond, which explains the planar structure of these parts of the molecules, as well as the molecules of dienes and arenes. sp 2 -Hybridization is also observed at carbon atoms and in such an allotropic modification of carbon as graphite.

In the molecules of some beryllium compounds, sp hybridization of the valence orbitals of the beryllium atom in the excited state is observed.

Two hybrid clouds are oriented relative to each other at an angle of 180° (Fig. 15), and therefore the beryllium chloride BeCl 2 molecule has a linear shape.

Rice. 15.
Schemes of sp-hybridization and location of sp-clouds in space

A similar type of hybridization of atomic orbitals exists for carbon atoms in alkynes - hydrocarbons of the acetylene series - at the site of the triple bond.

This hybridization of orbitals is characteristic of carbon atoms in another of its allotropic modifications - carbyne:

Table 9 shows the types of geometric configurations of molecules corresponding to certain types of hybridization of the orbitals of the central atom A, taking into account the influence of the number of free (non-bonding) electron pairs.

Table 9
Geometric configurations of molecules corresponding various types hybridization of outer electron orbitals of the central atom

Questions and tasks for § 7

In the molecules of hydrogen compounds of carbon, nitrogen and oxygen, the formulas of which are CH 4, NH 3 and H 2 O, the valence orbitals of the central non-metal atoms are in a state of sp 3 hybridization, but the bond angles between the bonds are different - 109°28" 107°30" and 104°27" respectively. How can this be explained?
Why is graphite electrically conductive and diamond not?
What geometric shape will the molecules of two fluorides have - boron and nitrogen (BF 3 and NF 3, respectively)? Give a reasoned answer.
The silicon fluoride molecule SiF 4 has a tetrahedral structure, and the bromine chloride molecule BCl 3 has a triangle shape - planar. Why?