Electronic configuration of the ground state of the carbon atom l s 2 2s 2 2p 2:

It would be expected that such a carbon atom would form a CH 2 compound with two hydrogen atoms. But in methane, the carbon is bonded to four hydrogen atoms. In order to represent the formation of a CH 4 molecule, it is necessary to refer to its excited electronic state.

Now one would expect the carbon atom to form four bonds with hydrogen atoms, but these bonds are not equivalent: three bonds are formed using R-orbitals, one - using s-orbitals. This contradicts the well-known fact that all bonds in methane are equivalent.

Next, hybridization of the orbitals is carried out. Mathematically calculated various combinations of one s- and three R-orbitals. The hybrid orbitals with the highest degree of directivity ("better" orbitals), as a result of more overlap, give bonds (1) stronger than unhybridized s- or R-orbitals. Four "best" hybrid orbitals (2) are equivalent . They are directed to the vertices of a regular tetrahedron, the angle between the two orbitals is 109.5 o. This geometry provides (3) minimal repulsion between them .

Let's complete the picture of the construction of the methane molecule: each of the four sp 3 -orbitals of a carbon atom overlaps with 1 s-orbital of the hydrogen atom. Four  -connections.

For maximum coverage sp 3 -orbitals of carbon and 1 s- orbitals of hydrogen four hydrogen atoms must lie on the axes sp 3 -orbitals. Therefore, the angle between any two C–H bonds is 109.5 o.

Experimental data confirm the calculation: methane has a tetrahedral structure. All carbon-hydrogen bonds have the same length of 10.9 10 -2 nm, the angle between any two bonds is tetrahedral and equal to 109.5 o. It takes 427·10 3 J/mol to break one of the bonds in methane.

1.3. The structure of ethane

The construction of the next homologue of the alkane series - ethane H 3 C–CH 3 will be carried out in a similar way. As in the case of methane , C-H bonds arise due to overlap sp 3 -orbitals of the carbon atom with 1s-orbitals of hydrogen atoms, the carbon-carbon bond is formed as a result of the overlap of two sp 3 -orbitals of carbon atoms.

The ethane molecule contains six carbon-hydrogen  bonds and one carbon-carbon  bond. -bonds have cylindrical symmetry  . The symmetry axis of the -bond electron cloud is the line connecting the atoms. The electron cloud of a carbon-carbon  bond, which has cylindrical symmetry, can be depicted as follows:

1.4. Rotation around a simple carbon-carbon bond. Conformations

In the ethane molecule, the rotation of one methyl group relative to the other occurs almost freely.

	Various arrangements of groups and atoms in space, resulting from the rotation of one atom relative to another along the bond line connecting these atoms, are calledconformations .
Shielded ethane(I) conformation		Hindered ethane(II) conformation

However, the rotation of one methyl group relative to another is not entirely free. The potential energy of the molecule is minimal for the hindered conformation II; during the rotation of the methyl group, it increases and reaches a maximum for the hindered conformation I. The excess energy of the hindered conformation compared to the hindered conformation is called the energy torsion stress . For an ethane molecule, this energy is 13 10 3 J/mol (Fig. 1.1).

It is believed that the excess energy appears due to the repulsion of electron clouds of carbon-hydrogen bonds at the moment when they pass by each other. Since at room temperature the energy of some collisions of molecules can reach 86·10 3 J/mol, the barrier of 13·10 3 J/mol is easily overcome. Rotation in ethane is considered to be free.

Rice. 1.1. Energyprofilegroup rotationsCH 3 in an ethane molecule around a carbon-carbon bond

Conformations corresponding to energy minima are called conformational isomers or conformers . In more complex molecules, the number of conformers can be quite large.

Topics of the USE codifier: Covalent chemical bond, its varieties and mechanisms of formation. Characteristics of a covalent bond (polarity and bond energy). Ionic bond. Metal connection. hydrogen bond

Intramolecular chemical bonds

Let us first consider the bonds that arise between particles within molecules. Such connections are called intramolecular.

chemical bond between atoms of chemical elements has an electrostatic nature and is formed due to interactions of external (valence) electrons, in more or less degree held by positively charged nuclei bonded atoms.

The key concept here is ELECTRONEGNATIVITY. It is she who determines the type of chemical bond between atoms and the properties of this bond.

is the ability of an atom to attract (hold) external(valence) electrons. Electronegativity is determined by the degree of attraction of external electrons to the nucleus and depends mainly on the radius of the atom and the charge of the nucleus.

Electronegativity is difficult to determine unambiguously. L. Pauling compiled a table of relative electronegativity (based on the bond energies of diatomic molecules). The most electronegative element is fluorine with meaning 4 .

It is important to note that in different sources you can find different scales and tables of electronegativity values. This should not be frightened, since the formation of a chemical bond plays a role atoms, and it is approximately the same in any system.

If one of the atoms in the chemical bond A:B attracts electrons more strongly, then the electron pair is shifted towards it. The more electronegativity difference atoms, the more the electron pair is displaced.

If the electronegativity values ​​of the interacting atoms are equal or approximately equal: EO(A)≈EO(V), then the shared electron pair is not displaced to any of the atoms: A: B. Such a connection is called covalent non-polar.

If the electronegativity of the interacting atoms differ, but not much (the difference in electronegativity is approximately from 0.4 to 2: 0,4<ΔЭО<2 ), then the electron pair is shifted to one of the atoms. Such a connection is called covalent polar .

If the electronegativity of the interacting atoms differ significantly (the difference in electronegativity is greater than 2: ΔEO>2), then one of the electrons almost completely passes to another atom, with the formation ions. Such a connection is called ionic.

The main types of chemical bonds are − covalent, ionic and metallic connections. Let's consider them in more detail.

covalent chemical bond

covalent bond – it's a chemical bond formed by formation of a common electron pair A:B . In this case, two atoms overlap atomic orbitals. A covalent bond is formed by the interaction of atoms with a small difference in electronegativity (as a rule, between two non-metals) or atoms of one element.

Basic properties of covalent bonds

• orientation,
• saturability,
• polarity,
• polarizability.

These bond properties affect the chemical and physical properties of substances.

Direction of communication characterizes the chemical structure and form of substances. The angles between two bonds are called bond angles. For example, in a water molecule, the H-O-H bond angle is 104.45 o, so the water molecule is polar, and in the methane molecule, the H-C-H bond angle is 108 o 28 ′.

Saturability is the ability of atoms to form a limited number of covalent chemical bonds. The number of bonds that an atom can form is called.

Polarity bonds arise due to the uneven distribution of electron density between two atoms with different electronegativity. Covalent bonds are divided into polar and non-polar.

Polarizability connections are the ability of bond electrons to be displaced by an external electric field(in particular, the electric field of another particle). The polarizability depends on the electron mobility. The farther the electron is from the nucleus, the more mobile it is, and, accordingly, the molecule is more polarizable.

Covalent non-polar chemical bond

There are 2 types of covalent bonding - POLAR and NON-POLAR .

Example . Consider the structure of the hydrogen molecule H 2 . Each hydrogen atom carries 1 unpaired electron in its outer energy level. To display an atom, we use the Lewis structure - this is a diagram of the structure of the external energy level of an atom, when electrons are denoted by dots. Lewis point structure models are a good help when working with elements of the second period.

H. + . H=H:H

Thus, the hydrogen molecule has one common electron pair and one H–H chemical bond. This electron pair is not displaced to any of the hydrogen atoms, because the electronegativity of hydrogen atoms is the same. Such a connection is called covalent non-polar .

Covalent non-polar (symmetrical) bond - this is a covalent bond formed by atoms with equal electronegativity (as a rule, the same non-metals) and, therefore, with a uniform distribution of electron density between the nuclei of atoms.

The dipole moment of nonpolar bonds is 0.

Examples: H 2 (H-H), O 2 (O=O), S 8 .

Covalent polar chemical bond

covalent polar bond is a covalent bond that occurs between atoms with different electronegativity (usually, different non-metals) and is characterized displacement common electron pair to a more electronegative atom (polarization).

The electron density is shifted to a more electronegative atom - therefore, a partial negative charge (δ-) arises on it, and a partial positive charge arises on a less electronegative atom (δ+, delta +).

The greater the difference in the electronegativity of atoms, the higher polarity connections and even more dipole moment . Between neighboring molecules and charges opposite in sign, additional attractive forces act, which increases strength connections.

Bond polarity affects the physical and chemical properties of compounds. The reaction mechanisms and even the reactivity of neighboring bonds depend on the polarity of the bond. The polarity of a bond often determines polarity of the molecule and thus directly affects such physical properties as boiling point and melting point, solubility in polar solvents.

Examples: HCl, CO 2 , NH 3 .

Mechanisms for the formation of a covalent bond

A covalent chemical bond can occur by 2 mechanisms:

1. exchange mechanism the formation of a covalent chemical bond is when each particle provides one unpaired electron for the formation of a common electron pair:

BUT . + . B= A:B

2. The formation of a covalent bond is such a mechanism in which one of the particles provides an unshared electron pair, and the other particle provides a vacant orbital for this electron pair:

BUT: + B= A:B

In this case, one of the atoms provides an unshared electron pair ( donor), and the other atom provides a vacant orbital for this pair ( acceptor). As a result of the formation of a bond, both electron energy decreases, i.e. this is beneficial for the atoms.

A covalent bond formed by the donor-acceptor mechanism, is not different by properties from other covalent bonds formed by the exchange mechanism. The formation of a covalent bond by the donor-acceptor mechanism is typical for atoms either with a large number of electrons in the external energy level (electron donors), or vice versa, with a very small number of electrons (electron acceptors). The valence possibilities of atoms are considered in more detail in the corresponding.

A covalent bond is formed by the donor-acceptor mechanism:

- in a molecule carbon monoxide CO(the bond in the molecule is triple, 2 bonds are formed by the exchange mechanism, one by the donor-acceptor mechanism): C≡O;

- in ammonium ion NH 4 +, in ions organic amines, for example, in the methylammonium ion CH 3 -NH 2 + ;

- in complex compounds, a chemical bond between the central atom and groups of ligands, for example, in sodium tetrahydroxoaluminate Na the bond between aluminum and hydroxide ions;

- in nitric acid and its salts- nitrates: HNO 3 , NaNO 3 , in some other nitrogen compounds;

- in a molecule ozone O 3 .

Main characteristics of a covalent bond

A covalent bond, as a rule, is formed between the atoms of non-metals. The main characteristics of a covalent bond are length, energy, multiplicity and directivity.

Chemical bond multiplicity

Chemical bond multiplicity - This the number of shared electron pairs between two atoms in a compound. The multiplicity of the bond can be quite easily determined from the value of the atoms that form the molecule.

for example , in the hydrogen molecule H 2 the bond multiplicity is 1, because each hydrogen has only 1 unpaired electron in the outer energy level, therefore, one common electron pair is formed.

In the oxygen molecule O 2, the bond multiplicity is 2, because each atom has 2 unpaired electrons in its outer energy level: O=O.

In the nitrogen molecule N 2, the bond multiplicity is 3, because between each atom there are 3 unpaired electrons in the outer energy level, and the atoms form 3 common electron pairs N≡N.

Covalent bond length

Chemical bond length is the distance between the centers of the nuclei of atoms that form a bond. It is determined by experimental physical methods. The bond length can be estimated approximately, according to the additivity rule, according to which the bond length in the AB molecule is approximately equal to half the sum of the bond lengths in the A 2 and B 2 molecules:

The length of a chemical bond can be roughly estimated along the radii of atoms, forming a bond, or by the multiplicity of communication if the radii of the atoms are not very different.

With an increase in the radii of the atoms forming a bond, the bond length will increase.

for example

With an increase in the multiplicity of bonds between atoms (whose atomic radii do not differ, or differ slightly), the bond length will decrease.

for example . In the series: C–C, C=C, C≡C, the bond length decreases.

Bond energy

A measure of the strength of a chemical bond is the bond energy. Bond energy is determined by the energy required to break the bond and remove the atoms that form this bond to an infinite distance from each other.

The covalent bond is very durable. Its energy ranges from several tens to several hundreds of kJ/mol. The greater the bond energy, the greater the bond strength, and vice versa.

The strength of a chemical bond depends on the bond length, bond polarity, and bond multiplicity. The longer the chemical bond, the easier it is to break, and the lower the bond energy, the lower its strength. The shorter the chemical bond, the stronger it is, and the greater the bond energy.

for example, in the series of compounds HF, HCl, HBr from left to right the strength of the chemical bond decreases, because the length of the bond increases.

Ionic chemical bond

Ionic bond is a chemical bond based on electrostatic attraction of ions.

ions are formed in the process of accepting or giving away electrons by atoms. For example, the atoms of all metals weakly hold the electrons of the outer energy level. Therefore, metal atoms are characterized restorative properties the ability to donate electrons.

Example. The sodium atom contains 1 electron at the 3rd energy level. Easily giving it away, the sodium atom forms a much more stable Na + ion, with the electron configuration of the noble neon gas Ne. The sodium ion contains 11 protons and only 10 electrons, so the total charge of the ion is -10+11 = +1:

+11Na) 2 ) 8 ) 1 - 1e = +11 Na +) 2 ) 8

Example. The chlorine atom has 7 electrons in its outer energy level. To acquire the configuration of a stable inert argon atom Ar, chlorine needs to attach 1 electron. After the attachment of an electron, a stable chlorine ion is formed, consisting of electrons. The total charge of the ion is -1:

+17Cl) 2 ) 8 ) 7 + 1e = +17 Cl — ) 2 ) 8 ) 8

Note:

• The properties of ions are different from the properties of atoms!
• Stable ions can form not only atoms, but also groups of atoms. For example: ammonium ion NH 4 +, sulfate ion SO 4 2-, etc. Chemical bonds formed by such ions are also considered ionic;
• Ionic bonds are usually formed between metals and nonmetals(groups of non-metals);

The resulting ions are attracted due to electrical attraction: Na + Cl -, Na 2 + SO 4 2-.

Let us visually generalize difference between covalent and ionic bond types:

metal chemical bond

metal connection is the relationship that is formed relatively free electrons between metal ions forming a crystal lattice.

The atoms of metals on the outer energy level usually have one to three electrons. The radii of metal atoms, as a rule, are large - therefore, metal atoms, unlike non-metals, quite easily donate outer electrons, i.e. are strong reducing agents

Intermolecular interactions

Separately, it is worth considering the interactions that occur between individual molecules in a substance - intermolecular interactions . Intermolecular interactions are a type of interaction between neutral atoms in which new covalent bonds do not appear. The forces of interaction between molecules were discovered by van der Waals in 1869 and named after him. Van dar Waals forces. Van der Waals forces are divided into orientation, induction and dispersion . The energy of intermolecular interactions is much less than the energy of a chemical bond.

Orientation forces of attraction arise between polar molecules (dipole-dipole interaction). These forces arise between polar molecules. Inductive interactions is the interaction between a polar molecule and a non-polar one. A non-polar molecule is polarized due to the action of a polar one, which generates an additional electrostatic attraction.

A special type of intermolecular interaction is hydrogen bonds. - these are intermolecular (or intramolecular) chemical bonds that arise between molecules in which there are strongly polar covalent bonds - H-F, H-O or H-N. If there are such bonds in the molecule, then between the molecules there will be additional forces of attraction .

Mechanism of education The hydrogen bond is partly electrostatic and partly donor-acceptor. In this case, an atom of a strongly electronegative element (F, O, N) acts as an electron pair donor, and hydrogen atoms connected to these atoms act as an acceptor. Hydrogen bonds are characterized orientation in space and saturation .

The hydrogen bond can be denoted by dots: H ··· O. The greater the electronegativity of an atom connected to hydrogen, and the smaller its size, the stronger the hydrogen bond. It is primarily characteristic of compounds fluorine with hydrogen , as well as to oxygen with hydrogen , less nitrogen with hydrogen .

Hydrogen bonds occur between the following substances:

— hydrogen fluoride HF(gas, solution of hydrogen fluoride in water - hydrofluoric acid), water H 2 O (steam, ice, liquid water):

— solution of ammonia and organic amines- between ammonia and water molecules;

— organic compounds in which O-H or N-H bonds: alcohols, carboxylic acids, amines, amino acids, phenols, aniline and its derivatives, proteins, solutions of carbohydrates - monosaccharides and disaccharides.

The hydrogen bond affects the physical and chemical properties of substances. Thus, the additional attraction between molecules makes it difficult for substances to boil. Substances with hydrogen bonds exhibit an abnormal increase in the boiling point.

for example As a rule, with an increase in molecular weight, an increase in the boiling point of substances is observed. However, in a number of substances H 2 O-H 2 S-H 2 Se-H 2 Te we do not observe a linear change in boiling points.

Namely, at boiling point of water is abnormally high - not less than -61 o C, as the straight line shows us, but much more, +100 o C. This anomaly is explained by the presence of hydrogen bonds between water molecules. Therefore, under normal conditions (0-20 o C), water is liquid by phase state.

Method of valence bonds (localized electron pairs) assumes that each pair of atoms in a molecule is held together by one or more shared electron pairs. Therefore, the chemical bond appears to be two-electron and two-center, i.e. located between two atoms. In the structural formulas of compounds, it is indicated by a dash:

H-Cl, H-H, H-O-H

Consider in the light Sun method, such features of communication as saturation, directivity and polarizability.

Valence atom - is determined by the number of unpaired (valence) electrons that can take part in the formation of a chemical bond. Valence is expressed in small integers and is equal to the number of covalent bonds. The valence of elements, which manifests itself in covalent compounds, is often called covalency. Some atoms have a variable valence, for example, carbon in the ground state has 2 unpaired electrons and will be two valent. When an atom is excited, it is possible to steam out the other two paired electrons and then the carbon atom will become four valent:

The excitation of an atom to a new valence state requires the expenditure of energy, which is compensated by the energy released during the formation of bonds.

Orientation of the covalent bond

Mutual overlapping of clouds can occur in different ways, due to their different shapes. Distinguish σ-, π- and δ-connections.

Sigma - connections are formed when clouds overlap along a line passing through the nuclei of atoms. Pi-bonds occur when clouds overlap on both sides of the line connecting the nuclei of atoms. Delta - communications are carried out when all four blades of d - electron clouds overlap, located in parallel planes.

σ– bond can occur when there is overlap along the line connecting the nuclei of atoms in the following orbitals: s— s -, s— R-, R– R-, d— d-orbitals, and d— s-, d— R- orbitals. σ– bond has the properties of a localized two-center bond, which it is.

π-bond can be formed by overlapping on both sides of the line connecting the nuclei of atoms of the following orbitals: R—R-, R—d-, d—d-, f—p-, f—d- and f—f- orbitals.

So, s- elements are capable of formation only σ– bonds, R- elements - σ– and π– bonds, d- elements - σ–, π– and δ-bonds, a f- elements - σ– , π– , δ-bonds. With the joint formation of π- and σ-bonds, a double bond is obtained. If two occur at the same time π-and σ-bond, a triple bond is formed. The number of bonds formed between atoms is called the bond multiplicity.

When establishing a connection with s— orbitals, due to their spherical shape, there is no preferential direction in space, for the most beneficial formation of covalent bonds. In the case R- orbitals, the electron density is unevenly distributed, so there is a certain direction in which the formation of a covalent bond is most likely.

Hybridization of atomic orbitals

Consider an example. Imagine that four hydrogen atoms are combined with a carbon atom and a methane molecule CH 4 is formed.

The picture shows what is happening, but does not explain how they behave s— and R- orbitals, in the formation of such compounds. Although R- the orbital has two parts turned relative to each other, but it can form only one bond. As a result, it can be assumed that in the methane molecule one hydrogen atom is attached to 2 s— orbitals of carbon, the rest - to 2 R- orbitals. Then, each hydrogen atom will be in relation to the other at an angle of 90 °, but this is not so. The electrons repel each other and diverge over a greater distance. What is actually happening?

As a result, all orbitals combine, rearrange and form 4 equivalent hybrid orbitals that are directed towards the vertices of the tetrahedron. Each of the hybrid orbitals contains a certain contribution 2 s— orbitals and some contributions 2 R- orbitals. Since 4 hybrid orbitals are formed by one 2 s— and three 2 R- orbitals, then this method of hybridization is called sp 3 -hybridization.

sp 3 hybridization of orbitals in a methane molecule

As can be seen from the figure, the configuration of hybrid orbitals allows four hydrogen atoms to form covalent bonds with a carbon atom, while the orbitals will be located relative to each other at an angle of 109.5 °.

The same type of hybridization is present in molecules such as NH 3 , H 2 O. On one of sp 3 - hybrid orbitals, in the NH 3 molecule, there is a lone electron pair, and the other three orbitals are used to connect with hydrogen atoms. In the H 2 O molecule, two hybrid orbitals of the oxygen atom are occupied by unshared electron pairs, while the other two are used for bonding with hydrogen atoms.

The number of hybrid orbitals is determined by the number of single bonds, as well as the number of unshared electron pairs in the molecule. These electrons are in hybrid orbitals. When the non-hybrid orbitals of two atoms overlap, a multiple bond is formed. For example, in an ethylene molecule, the bond is realized as follows:

sp 2 -hybridization of ethylene atoms

The planar arrangement of three bonds around each carbon atom suggests that in this case sp 2 -hybridization ( hybrid orbitals are formed by one 2 s— and two 2 R- orbitals ). At the same time, one 2 R- the orbital remains unused (non-hybrid). Orbitals will be located relative to each other at an angle of 120 °.

In the same way, a triple bond is formed in the acetylene molecule. In this case, it happens sp-hybridization atoms, i.e. hybrid orbitals are formed by one 2 s— and one 2 R- orbitals, and two 2 R Orbitals are non-hybrid. Orbitals are located relative to each other at an angle of 180 °

The following are examples of the geometric arrangement of hybrid orbitals.

Set of atomic orbitals	Set of hybrid orbitals	Geometric arrangement of hybrid orbitals	Examples
s,p	sp	Linear (angle 180°)	Be (CH 3) 2, HgCl 2 MgBr 2, CaH 2, BaF 2, C 2 H 2
s,p,p	sp 2	Planar trigonal (angle 120°)	BF 3, GaCl 3, InBr 3, TeI 3, C 2 H 4
s,p,p,p	sp 3	Tetrahedral (angle 109.5°)	CH 4, AsCl 4 -, TiCl 4, SiCl 4, GeF 4
s,p,p,d	sp2d	Flat square (90° angle)	Ni(CO) 4 , 2 -
s,p,p,p,d	sp 3 d	Trigonal bipyramidal (angles 120° and 90°)	PF 5 , PCl 5 , AsF 5
s,p,p,p,d,d	sp 3 d 2	Octahedral (90° angle)	SF 6 , Fe(CN) 6 3- , CoF 6 3-

Categories ,

carbon atom model

The valence electrons of a carbon atom are located in one 2s orbital and two 2p orbitals. 2p orbitals are located at an angle of 90° to each other, and the 2s orbital has spherical symmetry. Thus, the arrangement of carbon atomic orbitals in space does not explain the occurrence of bond angles 109.5°, 120°, and 180° in organic compounds.

To resolve this contradiction, the notion hybridization of atomic orbitals. To understand the nature of the three options for the arrangement of bonds of the carbon atom, ideas about three types of hybridization were needed.

We owe the emergence of the concept of hybridization to Linus Pauling, who did a lot to develop the theory of chemical bonding.

The concept of hybridization explains how a carbon atom changes its orbitals to form compounds. Below we will consider this process of orbital transformation step by step. At the same time, it should be borne in mind that the division of the hybridization process into stages or stages is, in fact, nothing more than a mental device that allows a more logical and accessible presentation of the concept. Nevertheless, the conclusions about the spatial orientation of the bonds of the carbon atom, which we will eventually come to, fully correspond to the real state of affairs.

Electronic configuration of the carbon atom in the ground and excited state

The figure on the left shows the electron configuration of a carbon atom. We are only interested in the fate of the valence electrons. As a result of the first step, which is called excitement or promotion, one of the two 2s electrons moves to a free 2p orbital. At the second stage, the hybridization process itself takes place, which can be somewhat conventionally imagined as a mixture of one s- and three p-orbitals and the formation of four new identical orbitals from them, each of which retains the properties of the s-orbital by one quarter and the properties of p-orbitals. These new orbitals are called sp 3 - hybrid. Here, the superscript 3 denotes not the number of electrons occupying the orbitals, but the number of p-orbitals that took part in the hybridization. Hybrid orbitals are directed to the vertices of the tetrahedron, in the center of which there is a carbon atom. Each sp 3 hybrid orbital contains one electron. These electrons participate in the third stage in the formation of bonds with four hydrogen atoms, forming bond angles of 109.5°.

sp3 - hybridization. methane molecule.

The formation of planar molecules with 120° bond angles is shown in the figure below. Here, as in the case of sp 3 hybridization, the first step is excitation. At the second stage, one 2s and two 2p orbitals participate in hybridization, forming three sp 2 -hybrid orbitals located in the same plane at an angle of 120° to each other.

Formation of three sp2 hybrid orbitals

One p-rorbital remains unhybridized and is located perpendicular to the plane of sp 2 hybrid orbitals. Then (third step) two sp 2 hybrid orbitals of two carbon atoms combine electrons to form a covalent bond. Such a bond, formed as a result of the overlap of two atomic orbitals along the line connecting the nuclei of an atom, is called σ-bond.

The formation of sigma and pi bonds in the ethylene molecule

The fourth stage is the formation of a second bond between two carbon atoms. The bond is formed as a result of the overlapping of the edges of unhybridized 2p orbitals facing each other and is called π-bond. The new molecular orbital is a set of two regions occupied by electrons of the π-bond - above and below the σ-bond. Both bonds (σ and π) together make up double bond between carbon atoms. And finally, the last, fifth step is the formation of bonds between carbon and hydrogen atoms using the electrons of the four remaining sp 2 hybrid orbitals.

Double bond in the ethylene molecule

The third and last type of hybridization is shown on the example of the simplest molecule containing a triple bond, the acetylene molecule. The first step is the excitation of the atom, the same as before. At the second stage, hybridization of one 2s and one 2p orbitals occurs with the formation of two sp-hybrid orbitals that are at an angle of 180°. And the two 2p orbitals necessary for the formation of two π bonds remain unchanged.

Formation of two sp-hybrid orbitals

The next step is the formation of a σ-bond between two sp-hybridized carbon atoms, then two π-bonds are formed. One σ bond and two π bonds between two carbons together make up triple bond. Finally, bonds are formed with two hydrogen atoms. The acetylene molecule has a linear structure, all four atoms lie on the same straight line.

We have shown how the three main types of molecular geometry in organic chemistry arise as a result of various transformations of the atomic orbitals of carbon.

Two methods can be proposed for determining the type of hybridization of various atoms in a molecule.

Method 1. The most general way, suitable for any molecules. Based on the dependence of the bond angle on hybridization:

a) bond angles of 109.5°, 107° and 105° indicate sp 3 hybridization;

b) a valence angle of about 120 ° - sp 2 - hybridization;

c) valence angle 180°-sp-hybridization.

Method 2. Suitable for most organic molecules. Since the type of bond (single, double, triple) is associated with geometry, it is possible to determine the type of its hybridization by the nature of the bonds of a given atom:

a) all bonds are simple - sp 3 -hybridization;

b) one double bond - sp 2 -hybridization;

c) one triple bond - sp-hybridization.

Hybridization is a mental operation of transforming ordinary (energetically most favorable) atomic orbitals into new orbitals, the geometry of which corresponds to the experimentally determined geometry of molecules.

I. Introduction. Stereochemical features of the carbon atom.

Stereochemistry is a part of chemistry devoted to the study of the spatial structure of molecules and the influence of this structure on the physical and chemical properties of a substance, on the direction and speed of their reactions. The objects of study in stereochemistry are mainly organic substances. The spatial structure of organic compounds is associated primarily with the stereochemical features of the carbon atom. These features depend, in turn, on the valence state (hybridization type).

In condition sp3- hybridization, the carbon atom is bonded to four substituents. If we imagine a carbon atom located in the center of a tetrahedron, then the substituents will be located at the corners of the tetrahedron. An example is the methane molecule, whose geometry is given below:

If all four substituents are the same (СH 4 , CCl 4), the molecule is a regular tetrahedron with valence angles 109 o 28". bonds - the tetrahedron becomes irregular.

In condition sp2- hybridization, the carbon atom is bonded to three substituents, with all four atoms lying in the same plane; bond angles are 120 o. Between two adjacent carbon atoms that are in the state sp2- hybridization, is established, as you know, not only the usual sigma -connection (when the maximum electron density is located exactly on an imaginary line connecting the nuclei of interacting atoms), but also a second bond of a special type. This so-called pi -connection formed by overlapping unhybridized R- orbitals.

The greatest overlap can be achieved with a parallel arrangement of p-orbitals: it is this position that is energetically more favorable, its violation requires the expenditure of energy to break the pi bond. Therefore, there is no free rotation around the carbon-carbon double bond (an important consequence of the lack of free rotation around the double bond is the presence of geometric isomers; see section II.2).

For a pi bond on a line connecting the nuclei of interacting atoms, the electron density is zero; it is maximal "above" and "under" the plane in which the connection between them lies. For this reason, the energy of a pi bond is less than that of a sigma bond, and in most organic reactions for compounds containing both pi and sigma bonds, less strong pi bonds are broken first.