What does sp3 hybridization mean. Hybridization of atomic orbitals and the geometry of molecules

sp3 hybridization

sp 3 -Hybridization - hybridization, in which atomic orbitals of one s- and three p-electrons (Fig. 1).

Rice. one. Education sp 3 hybrid orbitals

Four sp 3-hybrid orbitals are symmetrically oriented in space at an angle of 109°28" (Fig. 2).

Model of an atom with sp 3-hybrid orbitals

The spatial configuration of a molecule whose central atom is formed sp 3-hybrid orbitals - tetrahedron

Tetrahedral spatial configuration of a molecule whose central atom is formed sp 3-hybrid orbitals

hybridization atom orbital carbon

Examples of compounds for which sp 3-hybridization: NH 3 , POCl 3 , SO 2 F 2 , SOBr 2 , NH 4+ , H 3 O + . Also, sp 3-hybridization is observed in all saturated hydrocarbons (alkanes, cycloalkanes) and other organic compounds: CH 4, C 5 H 12, C 6 H 14, C 8 H 18, etc. The general formula of alkanes is: C n H 2n+2. The general formula of cycloalkanes is: C n H 2n. In saturated hydrocarbons, all chemical bonds are single, therefore, between the hybrid orbitals of these compounds, only at-overlapping.

Form a chemical bond, i.e. only unpaired electrons can create a common electron pair with a “foreign” electron from another atom. When writing electronic formulas, unpaired electrons are located one by one in the orbital cell.

atomic orbital is a function that describes the density of the electron cloud at each point in space around the nucleus of an atom. An electron cloud is a region of space in which an electron can be found with a high probability.

To harmonize the electronic structure of the carbon atom and the valency of this element, the concepts of excitation of the carbon atom are used. In the normal (unexcited) state, the carbon atom has two unpaired 2 R 2 electrons. In an excited state (when energy is absorbed) one of 2 s 2-electrons can pass to free R-orbital. Then four unpaired electrons appear in the carbon atom:

Recall that in the electronic formula of an atom (for example, for carbon 6 C - 1 s 2 2s 2 2p 2) large numbers in front of the letters - 1, 2 - indicate the number of the energy level. Letters s and R indicate the shape of the electron cloud (orbitals), and the numbers to the right above the letters indicate the number of electrons in a given orbital. All s- spherical orbitals

At the second energy level except 2 s-there are three orbitals 2 R-orbitals. These 2 R-orbitals have an ellipsoidal shape, similar to dumbbells, and are oriented in space at an angle of 90 ° to each other. 2 R-Orbitals denote 2 R X , 2R y and 2 R z according to the axes along which these orbitals are located.

Shape and orientation of p-electron orbitals

When chemical bonds are formed, the electron orbitals acquire the same shape. So, in saturated hydrocarbons, one s-orbital and three R-orbitals of a carbon atom to form four identical (hybrid) sp 3-orbitals:

It - sp 3 - hybridization.

Hybridization- alignment (mixing) of atomic orbitals ( s and R) with the formation of new atomic orbitals, called hybrid orbitals.

Four sp 3 -hybrid orbitals of carbon atom

Hybrid orbitals have an asymmetric shape, elongated towards the attached atom. Electron clouds repel each other and are located in space as far as possible from each other. At the same time, the axes of four sp 3-hybrid orbitals turn out to be directed to the vertices of the tetrahedron (regular triangular pyramid).

Accordingly, the angles between these orbitals are tetrahedral, equal to 109 ° 28".

The tops of electron orbitals can overlap with the orbitals of other atoms. If electron clouds overlap along a line connecting the centers of atoms, then such a covalent bond is called sigma () - bond. For example, in a C 2 H 6 ethane molecule, a chemical bond is formed between two carbon atoms by overlapping two hybrid orbitals. This is a connection. In addition, each of the carbon atoms with its three sp 3-orbitals overlap with s-orbitals of three hydrogen atoms, forming three -bonds.

Scheme of overlapping electron clouds in the ethane molecule

In total, three valence states with different types of hybridization are possible for a carbon atom. Except sp 3-hybridization exists sp 2 - and sp-hybridization.

sp 2 -Hybridization- mixing one s- and two R-orbitals. As a result, three hybrid sp 2 -orbitals. These sp 2 -orbitals are located in the same plane (with axes X, at) and are directed to the vertices of the triangle with an angle between the orbitals of 120°. unhybridized R-orbital is perpendicular to the plane of the three hybrid sp 2 orbitals (oriented along the axis z). Upper half R-orbitals are above the plane, the lower half is below the plane.

Type of sp 2-hybridization of carbon occurs in compounds with a double bond: C=C, C=O, C=N. Moreover, only one of the bonds between two atoms (for example, C=C) can be a bond. (The other bonding orbitals of the atom are directed in opposite directions.) The second bond is formed as a result of the overlap of non-hybrid R-orbitals on both sides of the line connecting the nuclei of atoms.

Orbitals (three sp 2 and one p) carbon atom in sp 2 - hybridization

Covalent bond formed by lateral overlap R-orbitals of neighboring carbon atoms is called pi()-bond.

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Due to less overlap of orbitals, the -bond is less strong than the -bond.

sp-Hybridization- this is mixing (alignment in form and energy) of one s- and one R-orbitals with the formation of two hybrid sp-orbitals. sp- Orbitals are located on the same line (at an angle of 180 °) and directed in opposite directions from the nucleus of the carbon atom. Two R-orbitals remain unhybridized. They are located mutually perpendicular to the directions of -bonds. On the image sp-orbitals are shown along the axis y, and the unhybridized two R-orbitals- along the axes X and z.

Atomic orbitals (two sp and two p) of carbon in the state of sp hybridization

The triple carbon-carbon bond CC consists of a -bond that occurs when overlapping sp-hybrid orbitals, and two -bonds.

The electronic structure of the carbon atom

Carbon, which is part of organic compounds, exhibits a constant valence. The last energy level of the carbon atom contains 4 electrons, two of which occupy the 2s orbital, which has a spherical shape, and two electrons occupy the 2p orbitals, which have a dumbbell shape. When excited, one electron from the 2s orbital can go to one of the vacant 2p orbitals. This transition requires some energy costs (403 kJ/mol). As a result, the excited carbon atom has 4 unpaired electrons and its electronic configuration is expressed by the formula 2s1 2p3 .

An excited carbon atom is able to form 4 covalent bonds due to 4 of its own unpaired electrons and 4 electrons of other atoms. So, in the case of the hydrocarbon methane (CH4), the carbon atom forms 4 bonds with the s-electrons of the hydrogen atoms. In this case, 1 bond should be formed type s-s(between the s-electron of a carbon atom and the s-electron of a hydrogen atom) and 3 p-s bonds (between 3 p-electrons of a carbon atom and 3 s-electrons of 3 hydrogen atoms). This leads to the conclusion that the four covalent bonds formed by the carbon atom are not equivalent. However, practical experience chemistry indicates that all 4 bonds in the methane molecule are absolutely equivalent, and the methane molecule has a tetrahedral structure with bond angles of 109 °, which could not be the case if the bonds were not equivalent. After all, only the orbitals of p-electrons are oriented in space along mutually perpendicular axes x, y, z, and the orbital of an s-electron has a spherical shape, so the direction of formation of a bond with this electron would be arbitrary. The theory of hybridization was able to explain this contradiction. L. Polling suggested that in any molecules there are no bonds isolated from each other. When bonds are formed, the orbitals of all valence electrons overlap. Several types of hybridization of electron orbitals are known. It is assumed that in the molecule of methane and other alkanes 4 electrons enter into hybridization.

Hybridization of carbon atom orbitals

Orbital hybridization is a change in the shape and energy of some electrons during the formation of a covalent bond, leading to more effective overlap of orbitals and increased bond strength. Hybridization of orbitals always occurs when electrons belonging to different types orbitals. 1. sp 3 -hybridization (the first valence state of carbon). With sp3 hybridization, 3 p-orbitals and one s-orbital of an excited carbon atom interact in such a way that orbitals are obtained that are absolutely identical in energy and symmetrically located in space. This transformation can be written like this:

s + px + py + pz = 4sp3

During hybridization, the total number of orbitals does not change, but only their energy and shape change. It is shown that the sp3 hybridization of the orbitals resembles a three-dimensional figure-eight, one of the blades of which is much larger than the other. Four hybrid orbitals are extended from the center to the tops regular tetrahedron at angles of 109.50. The bonds formed by hybrid electrons (for example, the s-sp 3 bond) are stronger than the bonds made by unhybridized p-electrons (for example, the s-p bond). because the hybrid sp3 orbital provides a larger area of electron orbital overlap than the unhybridized p orbital. Molecules in which sp3 hybridization is carried out have a tetrahedral structure. In addition to methane, these include methane homologues, inorganic molecules such as ammonia. The figures show a hybridized orbital and a tetrahedral methane molecule. Chemical bonds that arise in methane between carbon and hydrogen atoms are of type 2 y-bonds (sp3 -s-bond). Generally speaking, any sigma bond is characterized by the fact that the electron density of two interconnected atoms overlaps along the line connecting the centers (nuclei) of atoms. y-bonds correspond to the maximum possible degree of overlapping of atomic orbitals, so they are strong enough. 2. sp2 hybridization (second valence state of carbon). Occurs as a result of the overlap of one 2s and two 2p orbitals. The resulting sp2 hybrid orbitals are located in the same plane at an angle of 1200 to each other, and the unhybridized p orbital is perpendicular to it. The total number of orbitals does not change - there are four of them.

s + px + py + pz = 3sp2 + pz

The sp2 hybridization state occurs in alkene molecules, in carbonyl and carboxyl groups, i.e. in compounds containing a double bond. So, in the ethylene molecule, the hybridized electrons of the carbon atom form 3 y-bonds (two sp 2 -s type bonds between the carbon atom and hydrogen atoms and one sp 2 -sp 2 type bond between carbon atoms). The remaining unhybridized p-electron of one carbon atom forms a p-bond with the unhybridized p-electron of the second carbon atom. characteristic feature The p-bond is that the overlap of electron orbitals goes beyond the line connecting the two atoms. The overlap of orbitals goes above and below the y-bond connecting both carbon atoms. Thus, a double bond is a combination of y- and p-bonds. The first two figures show that in the ethylene molecule the bond angles between the atoms that form the ethylene molecule are 1200 (correspondingly, the orientations of the three sp2 hybrid orbitals in space). The third and fourth figures show the formation of a p-bond. ethylene (formation of y-bonds) ethylene (formation of pi-bonds) in chemical reactions. 3. sp-hybridization (the third valence state of carbon). In the state of sp-hybridization, the carbon atom has two sp-hybrid orbitals located linearly at an angle of 1800 to each other and two unhybridized p-orbitals located in two mutually perpendicular planes. sp- Hybridization is typical for alkynes and nitriles, i.e. for compounds containing a triple bond.

s + px + py + pz = 2sp + py + pz

So, in an acetylene molecule, the bond angles between atoms are 1800. Hybridized electrons of a carbon atom form 2 y-bonds (one sp-s bond between a carbon atom and a hydrogen atom and another sp-sp type bond between carbon atoms. Two unhybridized p-electrons of one carbon atom form two p-bonds with unhybridized p-electrons second carbon atom. The overlap of p-electron orbitals goes not only above and below the y-bond, but also in front and behind, and the total cloud of p-electrons has a cylindrical shape. Thus, a triple bond is a combination of one y-bond and two p-bonds The presence of less strong two p-bonds in the acetylene molecule ensures the ability of this substance to enter into addition reactions with the breaking of the triple bond.

Conclusion: sp3 hybridization is characteristic of carbon compounds. As a result of hybridization of one s-orbital and three p-orbitals, four hybrid sp3-orbitals are formed, directed to the vertices of the tetrahedron with an angle between the orbitals of 109°.

Most organic compounds have molecular structure. Atoms in substances with a molecular type of structure always form only covalent bonds with each other, which is also observed in the case of organic compounds. Recall that a covalent bond is a type of bond between atoms, which is realized due to the fact that atoms share a part of their outer electrons in order to acquire the electronic configuration of a noble gas.

By the number of socialized electron pairs, covalent bonds in organic matter ah can be divided into single, double and triple. These types of connections are indicated in the graphic formula, respectively, by one, two or three lines:

The multiplicity of the bond leads to a decrease in its length, so a single C-C connection has a length of 0.154 nm, double C=C bond - 0.134 nm, triple C≡C bond - 0.120 nm.

Types of bonds according to the way the orbitals overlap

As is known, orbitals can have different shapes, for example, s-orbitals are spherical, and p-dumbbell-shaped. For this reason, bonds can also differ in the way electron orbitals overlap:

ϭ-bonds - are formed when the orbitals overlap in such a way that the region of their overlap is intersected by a line connecting the nuclei. Examples of ϭ-bonds:

π-bonds - are formed when the orbitals overlap, in two areas - above and below the line connecting the nuclei of atoms. Examples of π bonds:

How to know when there are π- and ϭ-bonds in a molecule?

With a covalent type of bond, there is always a ϭ-bond between any two atoms, and it has a π-bond only in the case of multiple (double, triple) bonds. Wherein:

Single bond - always a ϭ-bond
A double bond always consists of one ϭ- and one π-bond
A triple bond is always formed by one ϭ and two π bonds.

Let us indicate these types of bonds in the propinoic acid molecule:

Hybridization of carbon atom orbitals

Orbital hybridization is the process by which orbitals that originally have different forms and the energies mix, forming in return the same number of hybrid orbitals, equal in shape and energy.

For example, when mixing one s- and three p- four orbitals are formed sp 3-hybrid orbitals:

In the case of carbon atoms, hybridization always takes part s- orbital, and the number p-orbitals that can take part in hybridization varies from one to three p- orbitals.

How to determine the type of hybridization of a carbon atom in an organic molecule?

Depending on how many other atoms a carbon atom is bonded to, it is either in the state sp 3, or in the state sp 2, or in the state sp- hybridization:

Let's practice determining the type of hybridization of carbon atoms using the example of the following organic molecule:

The first carbon atom is bonded to two other atoms (1H and 1C), so it is in the state sp-hybridization.

The second carbon atom is bonded to two atoms - sp-hybridization
The third carbon atom is bonded to four other atoms (two C and two H) - sp 3-hybridization
The fourth carbon atom is bonded to three other atoms (2O and 1C) - sp 2-hybridization.

Radical. Functional group

The term "radical" most often means a hydrocarbon radical, which is the remainder of a molecule of any hydrocarbon without one hydrogen atom.

The name of the hydrocarbon radical is formed based on the name of the corresponding hydrocarbon by replacing the suffix –en to suffix –silt .

Functional group - a structural fragment of an organic molecule (a certain group of atoms), which is responsible for its specific Chemical properties.

Depending on which of the functional groups in the molecule of the substance is the eldest, the compound is assigned to one or another class.

R is the designation of a hydrocarbon substituent (radical).

Radicals can contain multiple bonds, which can also be considered as functional groups, since multiple bonds contribute to the chemical properties of the substance.

If an organic molecule contains two or more functional groups, such compounds are called polyfunctional.

Concept of hybridization

The concept of hybridization of valence atomic orbitals was proposed by the American chemist Linus Pauling to answer the question why, if the central atom has different (s, p, d) valence orbitals, the bonds formed by it in polyatomic molecules with the same ligands are equivalent in their energy and spatial characteristics.

Ideas about hybridization are central to the method of valence bonds. Hybridization itself is not a real physical process, but only a convenient model that makes it possible to explain the electronic structure of molecules, in particular, hypothetical modifications of atomic orbitals during the formation of a covalent chemical bond, in particular, the alignment of chemical bond lengths and bond angles in a molecule.

The concept of hybridization was successfully applied to the qualitative description of simple molecules, but was later extended to more complex ones. Unlike the theory of molecular orbitals, it is not strictly quantitative, for example, it is not able to predict the photoelectron spectra of even such simple molecules as water. It is currently used mainly for methodological purposes and in synthetic organic chemistry.

This principle is reflected in the Gillespie-Nyholm theory of repulsion of electron pairs. The first and most important rule which was formulated as follows:

"Electronic pairs take such an arrangement on the valence shell of the atom, in which they are as far away from each other as possible, that is, electron pairs behave as if they repel each other."

The second rule is that "all electron pairs included in the valence electron shell are considered to be located at the same distance from the nucleus".

Types of hybridization

sp hybridization

Occurs when mixing one s- and one p-orbitals. Two equivalent sp-atomic orbitals are formed, located linearly at an angle of 180 degrees and directed towards different sides from the nucleus of a carbon atom. The two remaining non-hybrid p-orbitals are located in mutually perpendicular planes and participate in the formation of π-bonds, or are occupied by lone pairs of electrons.

sp 2 hybridization

Occurs when mixing one s- and two p-orbitals. Three hybrid orbitals are formed with axes located in the same plane and directed to the vertices of the triangle at an angle of 120 degrees. The non-hybrid p-atomic orbital is perpendicular to the plane and, as a rule, participates in the formation of π-bonds

sp 3 hybridization

Occurs when mixing one s- and three p-orbitals, forming four sp3-hybrid orbitals of equal shape and energy. They can form four σ-bonds with other atoms or be filled with lone pairs of electrons.

The axes of sp3-hybrid orbitals are directed to the vertices of a regular tetrahedron. The tetrahedral angle between them is 109°28", which corresponds to the lowest electron repulsion energy. Sp3 orbitals can also form four σ-bonds with other atoms or be filled with unshared pairs of electrons.

Hybridization and molecular geometry

Ideas about the hybridization of atomic orbitals underlie the Gillespie-Nyholm theory of repulsion of electron pairs. Each type of hybridization corresponds to a strictly defined spatial orientation of the hybrid orbitals of the central atom, which allows it to be used as the basis of stereochemical concepts in the world. organic chemistry.

The table shows examples of the correspondence between the most common types of hybridization and the geometric structure of molecules, assuming that all hybrid orbitals participate in the formation of chemical bonds (there are no unshared electron pairs).

Type of hybridization	Number hybrid orbitals	Geometry	Examples
sp	2	Linear	BeF 2 , CO 2 , NO 2 +
sp 2	3	triangular	BF 3, NO 3 -, CO 3 2-
sp 3	4	tetrahedral	CH 4, ClO 4 -, SO 4 2-, NH 4 +
dsp2	4	flat square	Ni(CO) 4 , XeF 4
sp 3 d	5	Hexahedral	PCl 5 , AsF 5
sp 3 d 2	6	Octahedral	SF 6 , Fe(CN) 6 3- , CoF 6 3-

Literature

Pauling L. Nature chemical bond/ Per. from English. M. E. Dyatkina. Ed. prof. Ya. K. Syrkina. - M.; L.: Goshimizdat, 1947. - 440 p.
Pauling L. general chemistry. Per. from English. - M .: Mir, 1974. - 846 p.
Minkin V. I., Simkin B. Ya., Minyaev R. M. Theory of the structure of molecules. - Rostov-on-Don: Phoenix, 1997. - S. 397-406. - ISBN 5-222-00106-7
Gillespie R. Geometry of molecules / Per. from English. E. Z. Zasorina and V. S. Mastryukov, ed. Yu. A. Pentina. - M .: Mir, 1975. - 278 p.

Notes

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In the process of determining the geometric shape of a chemical particle, it is important to take into account that pairs of valence electrons of the main atom, including those that do not form a chemical bond, are at a great distance from each other in space.

Term Features

When considering the issue of covalent chemical bonding, a concept is often used as the hybridization of atomic orbitals. This term is related to the alignment of form and energy. The hybridization of atomic orbitals is associated with the quantum-chemical process of rearrangement. Orbitals in comparison with the initial atoms have a different structure. The essence of hybridization lies in the fact that the electron that is located next to the nucleus of a bound atom is determined not by a specific atomic orbital, but by their combination with an equal principal quantum number. Basically, this process concerns higher, close in energy atomic orbitals that have electrons.

Process Specifics

The types of hybridization of atoms in molecules depend on how the orientation of new orbitals occurs. According to the type of hybridization, one can determine the geometry of an ion or molecule, suggest the features of chemical properties.

Types of hybridization

This type of hybridization, like sp, is a linear structure, the angle between the bonds is 180 degrees. An example of a molecule with a similar hybridization variant is BeCl 2 .

The next type of hybridization is sp 2 . Molecules are characterized by a triangular shape, the angle between bonds is 120 degrees. A typical example of such a hybridization variant is BCl 3 .

The sp 3 hybridization type suggests a tetrahedral structure of the molecule, a typical example of a substance with this hybridization variant is the methane CH 4 molecule. The bond angle in this case is 109 degrees 28 minutes.

Not only paired electrons, but also unseparated pairs of electrons are directly involved in hybridization.

Hybridization in a water molecule

For example, in a water molecule, there are two covalent polar bonds between the oxygen atom and the hydrogen atoms. In addition, the oxygen atom itself has two pairs of outer electrons that do not take part in the creation of a chemical bond. These 4 electron pairs in space occupy a certain place around the oxygen atom. Since they all have the same charge, they repel each other in space, the electron clouds are at a significant distance from each other. The type of hybridization of atoms in a given substance involves a change in the shape of atomic orbitals, they are stretched and aligned to the vertices of the tetrahedron. As a result, the water molecule acquires an angular shape, the bond angle between the oxygen-hydrogen bonds is 104.5 o.

To predict the type of hybridization, one can use the donor-acceptor mechanism of chemical bond formation. As a result, the free orbitals of an element with a lower electronegativity overlap, as well as the orbitals of an element with a higher electrical negativity, on which a pair of electrons is located. In the process of compiling the electronic configuration of an atom, their oxidation state is taken into account.

Rules for identifying the type of hybridization

In order to determine the type of carbon hybridization, certain rules can be used:

identify the central atom, calculate the number of σ-bonds;
put in the particle the oxidation state of atoms;
write down the electronic configuration of the main atom in the desired oxidation state;
make up the scheme of distribution along the orbits of valence electrons, pairing electrons;
allocate orbitals that are directly involved in the formation of bonds, find unpaired electrons (if the number of valence orbitals is insufficient for hybridization, orbitals of the next energy level are used).

The geometry of the molecule is determined by the type of hybridization. It is not affected by the presence of pi bonds. In the case of additional bonding, a change in the bond angle is possible, the reason is the mutual repulsion of electrons forming a multiple bond. So, in the nitric oxide molecule (4) during sp 2 hybridization, the bond angle increases from 120 degrees to 134 degrees.

Hybridization in the ammonia molecule

An unshared pair of electrons affects the resulting dipole moment of the entire molecule. Ammonia has a tetrahedral structure with an unshared pair of electrons. The ionicity of the nitrogen-hydrogen and nitrogen-fluorine bonds are 15 and 19 percent, the lengths are determined to be 101 and 137 pm, respectively. Thus, the nitrogen fluoride molecule should have a larger dipole moment, but the results of the experiment indicate the opposite.

Hybridization in organic compounds

Each class of hydrocarbons has its own type of hybridization. So, in the formation of molecules of the class of alkanes (saturated hydrocarbons), all four electrons of the carbon atom form hybrid orbitals. When they overlap, 4 hybrid clouds are formed, aligned to the vertices of the tetrahedron. Further, their tops overlap with non-hybrid s-orbitals of hydrogen, forming a single bond. Saturated hydrocarbons are characterized by sp 3 hybridization.

In unsaturated alkenes (their typical representative is ethylene), only three electron orbitals take part in hybridization - s and 2 p, three hybrid orbitals form a triangle in space. Non-hybrid p-orbitals overlap, creating a multiple bond in the molecule. This class of organic hydrocarbons is characterized by the sp 2 hybrid state of the carbon atom.

Alkynes differ from the previous class of hydrocarbons in that only two types of orbitals participate in the hybridization process: s and p. The two non-hybrid p-electrons remaining at each carbon atom overlap in two directions, forming two multiple bonds. This class of hydrocarbons is characterized by the sp-hybrid state of the carbon atom.

Conclusion

By determining the type of hybridization in a molecule, it is possible to explain the structure of various inorganic and organic substances, to predict the possible chemical properties of a particular substance.

An important characteristic of 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.

To explain the geometric configuration of the molecule, the concept of hybridization of the AO of the central atom is used. The excited beryllium atom has the 2s 1 2p 1 configuration, the excited boron atom has the 2s 1 2p 2 configuration, and the excited carbon atom has the 2s 1 2p 3 configuration. 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 , BCl 3 , CCl 4 should be unequal in energy and direction of bond. However, experimental data show that in molecules containing central atoms with different valence orbitals

(s, p, d), all connections are equivalent. To resolve this contradiction, Pauling and Slater proposed the concept of hybridization

The main provisions of the concept of hybridization:

1. Hybrid orbitals are formed from different atomic orbitals, not very different in energy,

2. The number of hybrid orbitals is equal to the number of atomic orbitals involved in hybridization.

3. Hybrid orbitals are the same in the shape of the electron cloud and in energy.

4 Compared to atomic orbitals, they are more elongated in the direction of formation of chemical bonds and therefore cause better overlap of electron clouds.

It should be noted that the hybridization of orbitals does not exist as a physical process. The hybridization method is a convenient model for the visual description of molecules.

Sp hybridization

sp–hybridization takes place, for example, in 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 corresponding energy level. When a molecule is formed, one s- and one p-orbital form two hybrid sp-orbitals at an angle of 180 o (Fig. 5).

Fig.5 sp hybrid orbitals

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

sp 2 hybridization

As a result of the combination 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° to each other. This is, for example, the configuration of the BF 3 molecule (Fig. 6):

Fig.6 sp 2 hybrid orbitals

sp 3 hybridization

sp 3 -Hybridization is characteristic of carbon compounds. As a result of the combination of one s-orbital and three p-orbitals, four hybrid sp 3 orbitals are formed, directed to 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 the carbon atom with other atoms in compounds, for example, in CH 4, CCl 4, C (CH 3) 4, etc. (Fig. 7).

Fig.7 sp 3 hybrid orbitals

The hybridization method explains the geometry of the ammonia molecule. As a result of the combination of one 2s and three 2p nitrogen orbitals, 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, and the fourth with a pair of electrons does not. angles between N-H bonds not equal to 90 o as in a pyramid, but not equal to 109.5 o corresponding to a tetrahedron (Fig. 8):

Fig.8 sp 3 - hybridization in the ammonia molecule

When ammonia interacts with a hydrogen ion H + + ׃NH 3 \u003d NH 4 +, 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 the angle between the O–H bonds in the corner water molecule. As a result of the combination of one 2s and three 2p oxygen orbitals, four sp 3 hybrid orbitals are formed, of which only two participate in the formation of a chemical bond, which leads to a distortion of the angle corresponding to the tetrahedron (Fig. 9):

Fig 9 sp 3 - hybridization in water molecule

Hybridization can include not only s- and p-, 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- (Fig. 10). In this case, the molecule has the configuration of an octahedron:

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

Ideas about hybridization make it possible to understand such features of the structure of molecules that cannot be explained in any other way. The hybridization of atomic orbitals (AO) leads to a shift of the electron cloud in the direction of bond formation with other atoms. As a result, the overlapping regions of hybrid orbitals turn out to be larger than for pure orbitals, and the bond strength increases.

Delocalized π-bond

According to the MVS method, the electronic structure of a molecule looks like a set of different valence schemes (localized pair method). But, as it turned out, it is impossible to explain the experimental data on the structure of many molecules and ions in terms of localized bonding only. Studies show that only σ-bonds are always localized. In the presence of π-bonds, there can be delocalization, at which the bonding electron pair simultaneously belongs to more than two atomic nuclei. For example, it has been experimentally established that the BF 3 molecule has a flat triangular shape (Fig. 6). All three links

B–F are equivalent, however, the value of the internuclear distance indicates that the bond is intermediate between single and double. These facts can be explained as follows. At the boron atom, as a result of the combination 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, but the free unhybridized p-orbital remains unused, and fluorine atoms have unshared electronic couples. Therefore, it is possible to form a π-bond by the donor-acceptor mechanism. The equivalence of all bonds indicates the delocalization of the π-bond between three fluorine atoms.

The structural formula of the BF 3 molecule, taking into account the delocalization of the π-bond, can be depicted as follows (the non-localized bond is indicated by a dotted line):

Rice.11 The structure of the BF 3 molecule

A non-localized π-bond determines the non-integer multiplicity of the bond. In this case, it is equal to 1 1/3 since between the boron atom and each of the fluorine atoms there is one σ-bond and 1/3 part of the π-bond.

In the same way, the equivalence of all bonds in the NO 3 - ion indicates the delocalization of the π-bond and the negative charge to all oxygen atoms. In a flat triangular ion NO 3 - (sp 2 -hybridization of the nitrogen atom) delocalized

π-bonds (depicted by dotted lines) are evenly distributed between all oxygen atoms (Fig. 12)

Rice. 12 Structural formula of the NO 3 ion - taking into account the delocalization of the π-bond

Similarly, delocalized π-bonds are evenly distributed between all oxygen atoms in anions: PO 4 3- (sp 3 - hybridization of the phosphorus atom → tetrahedron), SO 4 2- (sp 3 - hybridization of the sulfur atom → tetrahedron) (Fig. 13)