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The theory of valence played an important role in the development of the theory of chemistry in general and organic chemistry in particular. Proceeding from the theory of valence, Kekulé assumed that the carbon atom is tetravalent, and in 1858 tried, relying on this assumption, to represent the structure of the simplest organic molecules and radicals. In the same 1858, the Scottish chemist Archibald Scott Cooper (1831-1892) proposed depicting the forces connecting atoms (or bonds, as they are commonly called) in the form of dashes. After the first organic molecule was "built", it became quite clear why organic molecules, as a rule, are much larger and more complex than inorganic ones.

According to Kekulé's ideas, carbon atoms can bond to each other using one or more of their four valence bonds, forming long chains - straight or branched. Apparently, no other atoms have this remarkable ability to the extent that carbon does.

So, imagining that each carbon atom has four valence bonds, and each hydrogen atom has one such bond, we can depict the three simplest hydrocarbons (compounds whose molecules are formed only by carbon and hydrogen atoms), methane CH4, ethane C2H6 and propane C3H8, in the following way:

By increasing the number of carbon atoms, this sequence can be continued, and almost indefinitely. By adding oxygen (two valence bonds) or nitrogen (three valence bonds) to the hydrocarbon chain, we can represent the structural formulas of ethyl alcohol (C2H6O) and methylamine (CH5N) molecules:

Assuming the possibility of two bonds (double bond) or three bonds (triple bond) between neighboring atoms, one can depict the structural formulas of such compounds as ethylene (C2H4), acetylene (C2H2), methyl cyanide (C2H3N), acetone (C3H6O) and acetic acid (C2H4O2):

The usefulness of structural formulas was so obvious that many organic chemists adopted them immediately. They considered completely obsolete all attempts to portray organic molecules as structures built from radicals. As a result, it was considered necessary, while writing down the formula of a compound, to show its atomic structure.

The Russian chemist Alexander Mikhailovich Butlerov (1823-1886) used this new system of structural formulas in his theory of the structure of organic compounds. In the 60s of the last century, he showed how, with the help of structural formulas, one can clearly explain the reasons for the existence of isomers (see Chapter 5). So, for example, ethyl alcohol and dimethyl ether have the same empirical formula C2H6O, however, the structural formulas of these compounds differ significantly:

therefore, it is not surprising that a change in the arrangement of atoms results in two sets of very different properties. In ethyl alcohol, one in six hydrogen atoms is attached to an oxygen atom, while in dimethyl ether, all six hydrogen atoms are attached to carbon atoms. The oxygen atom holds the hydrogen atom weaker than the carbon atom, so that metallic sodium added to ethyl alcohol replaces hydrogen (one-sixth of the total). Sodium added to dimethyl ether does not displace hydrogen at all. Thus, when drawing up structural formulas, one can be guided by chemical reactions, and structural formulas, in turn, can help to understand the essence of the reactions.

Butlerov paid particular attention to one of the types of isomerism, called tautomerism (dynamic isomerism), in which some substances always act as mixtures of two compounds. If one of these compounds is isolated in pure form, it will immediately partially pass into the other compound. Butlerov showed that tautomerism is due to the spontaneous transition of a hydrogen atom from an oxygen atom to a neighboring carbon atom (and vice versa).

To fully prove the validity of the system of structural formulas, it was necessary to determine the structural formula of benzene - a hydrocarbon containing six carbon atoms and six hydrogen atoms. This was not done immediately. It seemed that there was no such structural formula that would, while meeting the requirements of valence, at the same time, would explain the greater stability of the compound. The first variants of the structural formulas of benzene were very similar to the formulas of some hydrocarbons - compounds that are very unstable and not similar in chemical properties to benzene.

Organic chemistry is the chemistry of the carbon atom. The number of organic compounds is tens of times greater than inorganic, which can only be explained features of the carbon atom :

a) he is in the middle of the electronegativity scale and the second period, therefore, it is not profitable for him to give up his own and accept other people's electrons and acquire a positive or negative charge;

b) special structure of the electron shell - there are no electron pairs and free orbitals (there is only one more atom with a similar structure - hydrogen, which is probably why carbon with hydrogen forms so many compounds - hydrocarbons).

The electronic structure of the carbon atom

С - 1s 2 2s 2 2p 2 or 1s 2 2s 2 2p x 1 2p y 1 2p z 0

Graphically:

A carbon atom in an excited state has the following electronic formula:

* С - 1s 2 2s 1 2p 3 or 1s 2 2s 1 2p x 1 2p y 1 2p z 1

As cells:

Shape of s- and p-orbitals

Atomic orbital - the region of space where it is most likely to detect an electron, with the corresponding quantum numbers.

It is a three-dimensional electronic "contour map" in which the wave function determines the relative probability of finding an electron at a given point in the orbital.

The relative sizes of atomic orbitals increase as their energies increase ( principal quantum number- n), and their shape and orientation in space is determined by the quantum numbers l and m. Electrons in orbitals are characterized by a spin quantum number. Each orbital can contain no more than 2 electrons with opposite spins.

When bonds are formed with other atoms, a carbon atom transforms its electron shell so that the strongest bonds are formed, and, therefore, as much energy as possible is released, and the system has acquired the greatest stability.

Changing the electron shell of an atom requires energy, which is then compensated by the formation of stronger bonds.

The transformation of the electron shell (hybridization) can be mainly of 3 types, depending on the number of atoms with which the carbon atom forms bonds.

Types of hybridization:

sp 3 - an atom forms bonds with 4 neighboring atoms (tetrahedral hybridization):

The electronic formula for sp 3 is a hybrid carbon atom:

* С –1s 2 2 (sp 3) 4 in the form of cells

The bond angle between hybrid orbitals is ~ 109 °.

Stereochemical formula of a carbon atom:

sp 2 - Hybridization (valence state)- an atom forms bonds with 3 neighboring atoms (trigonal hybridization):

Electronic formula of sp 2 - hybrid carbon atom:

* С –1s 2 2 (sp 2) 3 2p 1 in the form of cells

The bond angle between hybrid orbitals is ~ 120 °.

Stereochemical formula of sp 2 - hybrid carbon atom:

sp- Hybridization (valence state) - an atom forms bonds with 2 neighboring atoms (linear hybridization):

Electronic formula of sp - hybrid carbon atom:

* С –1s 2 2 (sp) 2 2p 2 in the form of cells

The bond angle between hybrid orbitals is ~ 180 °.

Stereochemical formula:

The s-orbital is involved in all types of hybridization, since it has a minimum of energy.

The rearrangement of the electron cloud allows the formation of the strongest bonds and the minimum interaction of atoms in the resulting molecule. Wherein hybrid orbitals may not be identical, but bond angles may be different, for example CH 2 Cl 2 and СCl 4

2. Covalent bonds in carbon compounds

Covalent bonds, properties, methods and reasons for education - the school curriculum.

Let me just remind you:

1. Communication formation between atoms can be considered as a result of overlapping of their atomic orbitals, and the more effective it is (the greater the overlap integral), the stronger the bond.

According to the calculated data, the relative overlapping efficiencies of atomic orbitals S rel increase as follows:

Consequently, the use of hybrid orbitals, for example, sp 3 -orbitals of carbon in the formation of bonds with four hydrogen atoms, leads to the emergence of stronger bonds.

2. Covalent bonds in carbon compounds are formed in two ways:

A)If two atomic orbitals overlap along their principal axes, then the resulting bond is called - σ-bond.

Geometry. Thus, during the formation of bonds with hydrogen atoms in methane, four hybrid sp 3 ~ orbitals of the carbon atom overlap with the s-orbitals of four hydrogen atoms, forming four identical strong σ-bonds located at an angle of 109 ° 28 "to each other (standard tetrahedral angle) A similar strictly symmetric tetrahedral structure also arises, for example, during the formation of CCl 4; if the atoms forming bonds with carbon are not the same, for example, in the case of CH 2 C1 2, the spatial structure will slightly differ from the completely symmetric one, although in essence it remains tetrahedral ...

Σ-bond length between carbon atoms depends on the hybridization of atoms and decreases when going from sp 3 - hybridization to sp. This is due to the fact that the s-orbital is closer to the nucleus than the p-orbital, therefore, the larger its fraction in the hybrid orbital, the shorter it is, and, consequently, the resulting bond is shorter.

B) If two atomic p -orbitals, located parallel to each other, carry out lateral overlap above and below the plane where the atoms are located, the resulting bond is called - π (pi) - communication

Side overlap of atomic orbitals is less efficient than overlapping along the main axis, therefore π -connections are less strong than σ -connection. This is manifested, in particular, in the fact that the energy of a double carbon-carbon bond is less than twice the energy of a single bond. Thus, the C-C bond energy in ethane is 347 kJ / mol, while the C = C bond energy in ethene is only 598 kJ / mol, and not ~ 700 kJ / mol.

The degree of lateral overlap of two atomic 2p orbitals , and hence the strength π -bond is maximum if two carbon atoms and four bonded to them atoms are located strictly in one plane, i.e. if they coplanar , since only in this case the atomic 2p orbitals are exactly parallel to each other and therefore are capable of maximum overlap. Any deviation from the coplanar state due to rotation around σ - a bond connecting two carbon atoms will lead to a decrease in the degree of overlap and, accordingly, to a decrease in strength π -bonding, which thus helps to maintain the flatness of the molecule.

Rotation around a carbon-carbon double bond is impossible.

Distribution π -electrons above and below the plane of the molecule means the existence areas of negative charge ready for interaction with any electron-deficient reagents.

Atoms of oxygen, nitrogen, etc. also have different valence states (hybridizations), while their electron pairs can be in both hybrid and p-orbitals.

Carbon (C) is the sixth element of the periodic table with an atomic weight of 12. The element belongs to non-metals and has an isotope of 14 C. The structure of the carbon atom underlies all organic chemistry, since all organic substances include carbon molecules.

Carbon atom

Carbon position in Mendeleev's periodic table:

• sixth serial number;
• fourth group;
• second period.

Rice. 1. Position of carbon in the periodic table.

Based on the data from the table, we can conclude that the structure of the atom of the element carbon includes two shells, on which six electrons are located. The valence of carbon, which is part of organic substances, is constant and equal to IV. This means that there are four electrons on the outer electronic level, and two on the inner one.

Of the four electrons, two occupy a spherical 2s orbital, and the remaining two occupy a 2p dumbbell-shaped orbital. In an excited state, one electron from the 2s orbital is transferred to one of the 2p orbitals. When an electron passes from one orbital to another, energy is expended.

Thus, an excited carbon atom has four unpaired electrons. Its configuration can be expressed by the formula 2s 1 2p 3. This makes it possible to form four covalent bonds with other elements. For example, in a methane (CH 4) molecule, carbon forms bonds with four hydrogen atoms - one bond between the s-orbitals of hydrogen and carbon and three bonds between the p-orbitals of carbon and s-orbitals of hydrogen.

The diagram of the structure of the carbon atom can be represented as + 6C) 2) 4 or 1s 2 2s 2 2p 2.

Rice. 2. The structure of the carbon atom.

Physical properties

Carbon occurs naturally in the form of rocks. Several allotropic modifications of carbon are known:

• graphite;
• diamond;
• carbyne;
• coal;
• soot.

All these substances differ in the structure of the crystal lattice. The hardest substance, diamond, has a cubic form of carbon. At high temperatures, diamond turns into graphite with a hexagonal structure.

Rice. 3. Crystal lattices of graphite and diamond.

Chemical properties

The atomic structure of carbon and its ability to attach four atoms of another substance determine the chemical properties of the element. Carbon reacts with metals to form carbides:

• Ca + 2C → CaC 2;
• Cr + C → CrC;
• 3Fe + C → Fe 3 C.

Also reacts with metal oxides:

• 2ZnO + C → 2Zn + CO 2;
• PbO + C → Pb + CO;
• SnO 2 + 2C → Sn + 2CO.

At high temperatures, carbon reacts with non-metals, in particular with hydrogen, forming hydrocarbons:

C + 2H 2 → CH 4.

With oxygen, carbon forms carbon dioxide and carbon monoxide:

• C + O 2 → CO 2;
• 2С + О 2 → 2СО.

Carbon monoxide is also formed when it interacts with water.

The simplest organic compound is methane. Its molecule consists of five atoms - one carbon atom and four hydrogen atoms, evenly distributed in space around this central carbon atom. Here we are faced, first of all, with the most important postulate of organic chemistry - in all uncharged organic molecules, carbon is always tetravalent. Graphically, this is expressed in the fact that it must be combined with the chemical symbols of other elements or the same carbon by four dashes. In methane, all four hydrogen atoms are at the same distance from the carbon atom and as far apart as possible in space.

In a methane molecule, a carbon atom is in the center of a regular tetrahedron, and four hydrogen atoms are at its vertices.

This is what a methane molecule looks like taking into account the size of the atoms.

To build a model of a molecule, we take a tetrahedron, that is, a regular tetrahedron made up of equilateral triangles, and place a carbon atom in its center. Hydrogen atoms will be located along the vertices of the tetrahedron. Let's combine all the hydrogens with the central carbon atom. The angle α between two such lines is 109 degrees and 28 minutes.

So, we have built a methane model. But what are the dimensions of the molecules in reality? In recent decades, using physical research methods (we will talk about them later), it has been possible to accurately determine the interatomic distances in the molecules of organic compounds. In a methane molecule, the distance between the centers of a carbon atom and any hydrogen atom is 0.109 nm (1 nanometer, nm, is 10 -9 m). To visualize what a molecule looks like in space, they use the Stewart-Brigleb models, in which atoms are represented by balls of a certain radius.

Now let us ask ourselves the following question: what are the forces that bind the atoms in the molecule of the organic compound, why do the hydrogen atoms not break away from the carbon center?

A carbon atom consists of a positively charged nucleus (its charge is +6) and six electrons occupying different orbitals * around the nucleus, each of which corresponds to a certain energy level.

* (The orbital can be considered as the region of space in which the probability of encountering an electron is greatest.)

Two electrons occupy the lowest orbital closest to the nucleus. They interact most strongly with "their" nucleus and do not take part in the formation of chemical bonds. The other four electrons are a different matter. It is believed that in the so-called unexcited carbon atom, that is, in a separate atom that does not form any bonds with other atoms, these electrons are located as follows: two on the lower sublevel s and two on a higher sublevel R... In a somewhat simplified and schematic way, we can assume that a cloud that forms an electron located on s-sub-level, has the shape of a sphere. Clouds R-electrons look like volumetric eights, and these eights can be located in space along the axes x, y and z... Accordingly, each atom has three R-orbitals: p x, p y and p z. So, each orbital in an atom has a certain shape and is located in a special way in space.

In order to interact with other atoms, to form chemical bonds with them, a carbon atom must first of all go into a special, excited state. In this case, one electron jumps from s-orbitals on p-orbital. As a result, one electron occupies a spherical s-orbital, and the other three electrons form three orbitals-eight. However, this position is energetically disadvantageous for the atom. The lower energy of the atom corresponds to four identical orbitals, symmetrically located in space. Therefore, there is mixing, averaging, or, as they say, hybridization available orbitals, and the result is four new identical orbitals.

These hybrid orbitals are also similar to eights, but the eights are one-sided: the electron density is almost completely shifted to one side. Such hybridized orbitals are denoted sp 3(according to the number of electrons from different non-hybrid orbitals participating in their formation: one with s-orbitals and three - with R-orbital).

How does the methane molecule work? To each of the four hybrid orbitals directed from the carbon atom in different directions (or rather, to the corners of an imaginary tetrahedron that can be built around it), hydrogen atoms are suitable H. A hydrogen atom is a nucleus with a charge of +1 (for a light isotope of ordinary hydrogen - just a proton), and one electron occupying a spherical orbital around the proton. Clouds of "carbon" and "hydrogen" electrons overlap, and this means the formation of a chemical bond. The more the clouds of electrons of different atoms overlap, the stronger the bond. Now it becomes clear why hybridized orbitals are more profitable - after all, such a one-sided, bulging figure eight can overlap much more strongly with the hydrogen electron cloud than non-hybrid orbitals less extended in space. Note that this reasoning is somewhat arbitrary: a pure, so to speak, single and unexcited carbon atom does not really exist. Therefore, it makes no sense to discuss how all these orbital transformations, called hybridization, actually occur. However, for the convenience of describing chemical bonds by means of formulas and numbers, such conventions are useful. We will be convinced of this more than once.

How to get methane?

One of the simplest ways is to act on aluminum carbide with water:

However, aluminum carbide is too expensive a raw material for obtaining such a common, such a cheap product as methane, and there is no need to obtain it from other compounds - after all, natural gas consists of 85-98% methane.

Methane is one of the main building blocks from which organic compounds can be built. What are these compounds and how can they be obtained from methane?

Actually, methane is a relatively inert substance, and the set of chemical reactions that can be carried out with it is small.

Let's take a mixture of two gases - methane and chlorine and place it in a glass vessel. If this vessel is kept in the dark, then no reaction is observed. But let's try to illuminate the bottle with sunlight ..

A quantum of light interacts with a chlorine molecule, as a result, the molecule splits into two parts - two chlorine atoms:

The resulting atoms are much more active than molecules; they immediately attack methane molecules and capture hydrogen atoms. In this case, molecules of hydrogen chloride HCl and very unstable, very active particles, the so-called metal radicals ⋅CH 3, are formed:

The result is a chlorine atom already known to us (its future fate is not difficult to predict: it attacks new methane molecules, and everything repeats itself) and chloromethane, or methyl chloride, is a derivative of methane, in which one of the hydrogen atoms is replaced by chlorine.

The reaction that we have described belongs to the category of the so-called chain reactions, in which each stage, as in a chain, is connected with the previous and the next. Active particles - the product of one stage (here these are chlorine atoms and ⋅СН 3 methyl radicals) - are used in the next stage as starting materials. The discovery of chain reactions was one of the major events in the history of chemical science, and Academician N. N. Semenov and the English scientist S. N. Khishelwood were awarded the Nobel Prize for their contribution to the study of such reactions and the creation of their theory.

If we introduce such quantities of reagents into the reaction that there are two molecules of methane per chlorine molecule, then basically we will get methyl chloride CH 3 Cl. If we take chlorine in excess, then the substitution reaction will go further and, in addition to methyl chloride, we will get methylene chloride CH 2 Cl 2, chloroform CHCl 3 and, finally, the product of complete replacement of hydrogen with chlorine, carbon tetrachloride CCl 4:

But let's not forget about our task: to build various complex molecules from simple methane building blocks. For this we need methyl chloride. If you act on this compound with metallic sodium, then out of every two CH 3 Cl molecules, one ethane molecule is formed, in which there is a carbon-carbon bond:

What is ethane? This is methane, in which one of the hydrogens is replaced by the methyl ⋅CH 3 radical. And this radical itself, as we already know, is obtained by the separation of one hydrogen atom from methane.

If now in ethane we replace one of the hydrogens (any atom) with methyl, then we get a new substance - propane CH 3 -CH 2 -CH 3. How this can be done in practice, we know: first, in methane and ethane, replace one hydrogen with chlorine and then act on a mixture of methyl and ethyl chloride sodium (this reaction is called the Würz reaction in honor of the French chemist who discovered it):

Let's go further. Let's replace one of the hydrogen atoms with chlorine in propane. It turns out that now it doesn't matter which atom to replace! Replacing hydrogen at the extreme carbon atom (there are two such atoms) or at the average, we get two different compounds: normal propyl chloride ( n-propyl chloride) and isopropyl chloride:

Let us now replace the chlorine atoms in each of these compounds with methyl groups. We get two different butane - normal (i.e. not branched) butane ( n-butane) and iso-butane:

Let's add to the obtained molecules one more "brick". Let's start with n-butane. Here you can substitute methyl for one of the extreme hydrogen atoms. We get normal pentane. You can replace one of the medium hydrogens. Let's come to iso-pentane. Apparently from n-butane you won't get anything new anymore. Let's turn to iso-butane. If one of the extreme hydrogens (in CH 3 -groups) is replaced in it, then we come to the already mentioned iso-pentane, and replacing the middle single hydrogen atom, we get neopentane:

You can continue this procedure indefinitely. All these connections are called hydrocarbons(more precisely - saturated hydrocarbons or alkanes), because they consist of only two elements - carbon and hydrogen. In any alkane, the number of hydrogen atoms is 2 n+ 2, where n- the number of carbon atoms. Therefore, the formula for the limiting hydrocarbon can be written in general form as follows: C n N 2n + 2 .

In building our structures, I must say, we stopped in time. The fact is that the number of possible isomers increases catastrophically rapidly with an increase in the number of carbon atoms in the alkane molecule. So, for decane, a С 10 Н 22 hydrocarbon, possibly 75 different isomers, the number of isomers for a С 20 Н 42 hydrocarbon (eicosane) is 366 319. The number of possible isomers for tetracontane, a С 40 Н 82 hydrocarbon is hard to imagine: 62 491 178 805 831.

Now it becomes clear why such a huge number of organic compounds are already known today - several million - and why, in this respect, organic chemistry has far outstripped inorganic chemistry. But until now, only the simplest representatives of organic substances have been spoken about - about saturated hydrocarbons.

We have derived a number of isomeric hydrocarbons from methane using the Wurtz reaction. However, in practice, no one does this. The fact is that the simplest hydrocarbons, along with methane, are contained in natural gas, the composition of which is different for different fields. For example, the gas from the Severo-Stavropol field contains 85% methane, about 5% ethane, 2.5% propane, and 1.4% pentane and heavier hydrocarbons. Gas of the Gazlinskoye field consists of methane by 98%, ethane in it is only 1.6%. There are many hydrocarbons in oil, but more on that in the next chapters.

Lower hydrocarbons - methane, ethane, propane and butane - are colorless, odorless gases or with a faint smell of gasoline. Hydrocarbons from pentane to pentadecane C 15 H 32 are liquids and, finally, higher hydrocarbons at ordinary temperatures are solids.

As the number of carbon atoms increases, the boiling point and melting point of the compound increases.

Saturated hydrocarbons have another name - paraffins reflecting their chemical inertness (Latin parum affinis- low affinity). And yet they are widely used in the chemical industry to obtain a wide variety of substances. The main directions of industrial use of methane are shown in the diagram.

Before ending the conversation about methane and saturated hydrocarbons, let us answer one question: how is the bond in paraffins between two carbon atoms carried out, for example, in ethane? Everything is simple here - around each carbon atom there are, as in methane, four hybridized sp 3 -orbitals, three of them make bonds with hydrogen atoms, and one overlaps with exactly the same orbital of another carbon atom. The C-C bond length is 0.154 nm.

The carbon atom is the basic building block of which organic compounds are built. In order to fill the valence shell with eight electrons (like an inert gas), a carbon atom must pair its electrons with the electrons of four hydrogen atoms.

As a result of hybridization and pairing of electrons, both the valence electron shell of carbon and the valence shells of hydrogen atoms are filled. An extremely stable electronic configuration is created and a stable CH 4 molecule called methane is formed.

The electrons of different atoms form pairs, which are symbolically indicated by dots. Each pair of such electrons gives a covalent bond. For convenience, each of these pairs of electrons from different atoms, or a covalent bond, is usually depicted as a line (bond) connecting linked atoms.

The four bonds from the carbon atom represent the four valences that the carbon atom possesses. Likewise, one bond between each hydrogen atom and a carbon atom represents one valency that each hydrogen atom possesses.

However, these simplified representations do not reflect the true 3D geometry of the methane molecule. Methane has a tetrahedral structure due to hybridization. All angles between НСН bonds in methane are equal to 109.5 ° (Fig. 8). The tetrahedral structure allows each of the hydrogen atoms to occupy a position as distant as possible from the neighboring hydrogen atoms. As a result, the repulsive forces between neighboring hydrogen atoms become minimal.

The carbon-hydrogen covalent bonds in methane are strong bonds. To obtain from 1 g-molecule (1 mole) of methane (16 g) its constituent carbon and hydrogen atoms, it would take 404 kcal of energy. Since there are four carbon-hydrogen bonds in the methane molecule, each has an average energy of 101 kcal / mol. Such a bond is considered to be a very strong covalent bond.