The belonging of substances to the group of saturated hydrocarbons is determined by the nature of their structure. Consider the structure of the simplest hydrocarbon - methane.

Methane CH 4 is a colorless and odorless gas, almost twice as light as air. It is formed in nature as a result of decomposition without air access of the remains of plant and animal organisms. Therefore, it can be found, for example, in swampy reservoirs, in coal mines. Methane is found in significant quantities in natural gas, which is now widely used as a fuel in everyday life and in industry.

In the methane molecule, the chemical bonds of the hydrogen atoms with the carbon atom are covalent. If the electron clouds overlapping in pairs during the formation of bonds are denoted by two dots or a valence line, the structure of methane can be expressed by the formulas:

Or

When in organic chemistry the study of the spatial structure of molecules began to develop, it was found that the methane molecule actually has a tetrahedral shape, and not a flat one, as we depict on paper.

Let us find out why the methane molecule is a tetrahedron. We must start, obviously, from the structure of the carbon atom. But here we run into a contradiction. Carbon atoms have four valence electrons, two of them are paired s-electrons, they cannot form chemical bonds with hydrogen atoms. Chemical bonds can only be established by two unpaired p-electrons. But then the methane molecule should have the formula not CH 4, but CH 2, which is not true. This contradiction is eliminated by the following interpretation of the formation of chemical bonds.

When a carbon atom interacts with hydrogen atoms, the s-electrons of the outer layer in it are steamed out, one of them occupies the vacant place of the third p-electron and forms a cloud in the form of a volume eight, perpendicular to the clouds of the other two p-electrons . The atom then passes, as they say, into an excited state. Now all four valence electrons have become unpaired, they can form four chemical bonds. But a new contradiction arises. Three p-electrons must form three chemical bonds with hydrogen atoms in mutually perpendicular directions, i.e. at an angle of 90 °, and the fourth hydrogen atom could join in an arbitrary direction, since the s-electron cloud has a spherical shape and these bonds, Obviously, they would differ in properties. Meanwhile, it is known that all C-H bonds in the methane molecule are the same and are located at an angle of 109 ° 28 ". The concept of hybridization of electron clouds helps to resolve this contradiction.

In the process of formation of chemical bonds, the clouds of all valence electrons of the carbon atom (one s- and three p-electrons) align and become the same. At the same time, they take the form of asymmetric volume eights, elongated towards the vertices of the tetrahedron (an asymmetric distribution of electron density means that the probability of finding an electron on one side of the nucleus is greater than on the other).

The angle between the axes of hybrid electron clouds turns out to be equal to 109°28", which allows them, as similarly charged, to move away from each other as much as possible. Being elongated to the vertices of the tetrahedron, such clouds can significantly overlap with the electron clouds of hydrogen atoms, which leads to a greater release of energy and the formation of strong chemical bonds with identical properties (Fig. A).

DEFINITION

Methane- the simplest representative of the class of saturated hydrocarbons (the structure of the molecule is shown in Fig. 1). It is a colorless, light, flammable gas, odorless and almost insoluble in water.

Its boiling point is -161.5 o C, the solidification point is -182.5 o C. A mixture of methane with air is extremely explosive (especially in a ratio of 1:10).

Rice. 1. The structure of the methane molecule.

Getting methane

Methane is quite common in nature. It is the main component of the natural gas of gas fields (up to 97%), it is contained in significant quantities in associated petroleum gas (released during oil production), as well as in coke oven gas. It is emitted from the bottom of swamps, ponds and stagnant waters, where it is formed during the decomposition of plant residues without air access, which is why methane is also called swamp gas. Finally, methane constantly accumulates in coal mines, where it is called firedamp.

Synthetic methods for producing methane show the relationship of inorganic substances with organic ones. It is possible to distinguish industrial (1, 2, 3) and laboratory (4, 5) methods of its production:

C + 2H 2 → CH 4 (kat = Ni, t 0) (1);

CO + 3H 2 → CH 4 + H 2 O (kat = Ni, t = 200 - 300 o C) (2);

CO 2 + 4H 2 → CH 4 + 2H 2 O (kat, t 0) (3);

Al 4 C 3 + 12H 2 O → CH 4 + 4Al(OH) 3 (4);

CH 3 COONa + NaOH → CH 4 + Na 2 CO 3 (5).

Chemical properties of methane

Methane is a low reactive organic compound. So, under normal conditions, it does not react with concentrated acids, molten and concentrated alkalis, alkali metals, halogens (except fluorine), potassium permanganate and potassium dichromate in an acidic medium.

All chemical transformations characteristic of methane proceed with the splitting of C-H bonds:

• halogenation (S R)

CH 4 + Cl 2 → CH 3 Cl + HCl ( hν);

• nitration (S R)

CH 4 + HONO 2 (dilute) → CH 3 -NO 2 + H 2 O (t 0);

• sulphochlorination (S R)

CH 4 + SO 2 + Cl 2 → CH 3 -SO 2 Cl + HCl ( hν);

There are catalytic (copper and manganese salts are used as catalysts) (1, 2, 3) and complete (combustion) (4) oxidation of methane:

2CH 4 + O 2 → 2CH 3 OH (p, t 0) (1);

CH 4 + O 2 → HC(O)H + H 2 O (NO, t 0) (2);

2CH 4 + 3O 2 → 2HCOOH + 2H 2 O (kat = Pt, t 0) (3);

CH 4 + 2O 2 → CO 2 + 2H 2 O + Q (4).

The conversion of methane with water vapor and carbon dioxide can also be attributed to the methods of its oxidation:

CH 4 + H 2 O →CO + 3H 2 (kat = Ni, t = 800 o C);

CH 4 + CO 2 → 2CO + 2H 2.

Methane cracking is the most important method of chemical processing of oil and its fractions in order to obtain products of lower molecular weight - lubricating oils, motor fuels, etc., as well as raw materials for the chemical and petrochemical industries:

2CH 4 → HC≡CH + 3H 2 (t = 1500 o C).

Methane application

Methane is the raw material basis of the most important chemical industrial processes for producing carbon and hydrogen, acetylene, oxygen-containing organic compounds- alcohols, aldehydes, acids.

Examples of problem solving

EXAMPLE 1

EXAMPLE 2

Exercise	Calculate the volumes of chlorine and methane, reduced to normal conditions, that will be required to obtain carbon tetrachloride with a mass of 38.5 g.
Solution	Let us write the equation for the reaction of methane chlorination to carbon tetrachloride (the reaction occurs under the action of UV radiation): CH 4 + 4Cl 2 \u003d CCl 4 + 4HCl. Calculate the amount of carbon tetrachloride substance ( molar mass equal to - 154 g / mol): n(CCl 4) \u003d m (CCl 4) / M (CCl 4); n (CCl 4) \u003d 38.5 / 154 \u003d 0.25 mol. According to the reaction equation n(CCl 4) : n(CH 4) = 1:1, i.e. n (CCl 4) \u003d n (CH 4) \u003d 0.25 mol. Then the volume of methane will be equal to: V(CH 4) = n(CH 4) × V m ; V (CH 4) \u003d 0.25 × 22.4 \u003d 5.6 l. According to the reaction equation, we find the amount of chlorine substance. n(CCl 4) : n(Cl 2) = 1:4, i.e. n(Cl 2) \u003d 4 × n (CCl 4) \u003d 4 × 0.25 \u003d 1 mol. Then the volume of chlorine will be equal to: V (Cl 2) \u003d n (Cl 2) × V m; V (Cl 2) \u003d 1 × 22.4 \u003d 22.4 l.
Answer	The volumes of chlorine and methane are 22.4 and 5.6 liters, respectively.

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. So the location atomic orbitals carbon 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 by 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.

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. That's why chemical bond appears to be two-electron and two-center, i.e. located between two atoms. AT structural formulas connections are 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 various 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 with the overlap of all four blades d - electron clouds located in parallel planes.

σ– bond may 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 ,

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).

Able 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.

Able 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 the pi bond is less than that of the sigma bond, and in most organic reactions for compounds containing both pi- and sigma-bonds, the weaker pi-bonds break first.