Organic chemistry is the chemistry of the carbon atom. The number of organic compounds is tens of times greater than inorganic ones, which can only be explained features of the carbon atom :
a) he is in middle of the electronegativity scale and the second period, therefore it is unprofitable for him to give 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 and hydrogen form so many compounds - hydrocarbons).
The electronic structure of the carbon atom
C - 1s 2 2s 2 2p 2 or 1s 2 2s 2 2p x 1 2p y 1 2p z 0
Graphically:
An excited carbon atom has the following electronic formula:
*C - 1s 2 2s 1 2p 3 or 1s 2 2s 1 2p x 1 2p y 1 2p z 1
In the form of cells: 
The shape of s- and p-orbitals
atomic orbital - the region of space where the electron is most likely to be found, 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 orbit.
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, the carbon atom transforms its electron shell so that the strongest bonds are formed, and, consequently, as much energy as possible is released, and the system acquires the greatest stability.
To change the electron shell of an atom, energy is required, 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 sp 3 - 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 carbon atom:
sp 2 – Hybridization (valence state)– an atom forms bonds with 3 neighboring atoms (trigonal hybridization):
The electronic formula 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 sp 2 - hybrid carbon atom:
sp– Hybridization (valence state) - the atom forms bonds with 2 neighboring atoms (linear hybridization):
The electronic formula of sp is a 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, because 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 the bond angles may be different, for example CH 2 Cl 2 and CCl 4
2. Covalent bonds in carbon compounds
Covalent bonds, properties, methods and causes of education - the school curriculum.
Let me just remind you:
1. Communication education between atoms can be considered as a result of the overlap 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 atomic orbital overlap efficiencies S rel increase as follows:
Therefore, the use of hybrid orbitals, for example, sp 3 orbitals of carbon in the formation of bonds with four hydrogen atoms, leads to stronger bonds.
2. Covalent bonds in carbon compounds are formed in two ways:
BUT)If two atomic orbitals overlap along their principal axes, then the resulting bond is called - σ bond.
Geometry. So, when bonds are formed with hydrogen atoms in methane, four hybrid sp 3 ~orbitals of a carbon atom overlap with 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 symmetrical tetrahedral structure also arises, for example, during the formation of CCl 4, but if the atoms that form bonds with carbon are not the same, for example in the case of CH 2 C1 2, the spatial structure will differ somewhat from completely symmetrical, although it remains essentially tetrahedral .
σ-bond length between carbon atoms depends on the hybridization of atoms and decreases in the transition from sp 3 - hybridization to sp. This is explained by the fact that the s orbital is closer to the nucleus than the p orbital, therefore, the larger its share in the hybrid orbital, the shorter it is, and therefore the shorter the resulting bond.
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, then the resulting bond is called - π (pi) - communication
Lateral overlap atomic orbitals is less efficient than overlapping along the principal axis, so π -bonds are less strong than σ -connections. This is manifested, in particular, in the fact that the energy of a double carbon-carbon bond exceeds the energy of a single bond by less than two times. 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.
Degree of lateral overlap of two atomic 2p orbitals , and hence the strength π -bond is maximum if two carbon atoms and four associated with them atoms are located strictly in the same plane, i.e. if they coplanar , since only in this case the atomic 2p orbitals are exactly parallel to each other and therefore capable of maximum overlap. Any deviation from coplanar due to rotation around σ -bond connecting two carbon atoms will lead to a decrease in the degree of overlap and, accordingly, to a decrease in strength π -bond, 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 to interact with any electron-deficient reagents.
The atoms of oxygen, nitrogen, etc. also have different valence states (hybridizations), while their electron pairs can be in both hybrid and p-orbitals.
In the ammonia molecule, the electrons around the nitrogen atom are also located in sp 3 hybridized orbitals. A similar picture is observed in the case of the water molecule.
NH3H2O 
With sp 3 -hybridization of orbitals, the carbon atom can give only simple s-bonds. When carbon forms a double bond, sp 2 hybridization is used (Fig. 7). In this case, one 2s and two 2p orbitals take part in hybridization, and one 2p orbital remains non-hybrid. The sp 2 orbitals are equivalent, their axes are coplanar and form an angle of 120° with each other; the non-hybrid 2p orbital is perpendicular to the plane of the hybrid orbitals.
Rice. 7 s and
two2 p orbitals to form three sp 2 hybrid orbitals.
When carbon forms a triple bond, sp hybridization is used. In this case, one 2s and one p orbitals take part in hybridization, and two 2p orbitals remain non-hybrid (Fig. 8).
Rice. eight Depiction of the mathematical hybridization procedure of one 2 s and one
noah2 p orbitals to form two sp hybrid orbitals.
acetylene
Exercise 13. Describe the bonds between atoms in molecules of (a) ethanoic acid, (b) ethanal, ethanamide in terms of atomic orbitals and predict all bond angles.
Answer(a) 
Bond length and energy
Shared electron pairs hold two bonding atoms at a certain distance called bond length. The bond length between atoms is approximately equal to the sum of their covalent radii (r) (Table 2), which makes it possible to calculate the lengths of any bonds. l A - B = r A + r B
table 2
Covalent radii (r) of some elements, Å
Ex. 14. In calculate the bond lengths for (a) C-H, (b) C-C, (c) C=C and (d) CºC,
(e) C-O, (f) C=O, (g) C-Cl, ... The values of covalent radii are given in Table. 1.2.
Answer(a) 0.77 + 0.37 = 1.14 Å, (b) 2 x 0.77 = 1.54 Å, (c) 2 x 0.67 = 1.34 Å, (d) ...
AT general case as the number of bonds between two atoms increases, their length decreases. In some molecules, the carbon-carbon bond length is intermediate between the length of single (1.54Å) and double (1.33Å) bonds. In this case, we talk about the order of communication. The approximate value of the order of such a relationship can be found graphically.
Energy is the ability to do work. A moving object has kinetic energy. If objects attract or repel each other, then they have potential energy. Two balls connected by a spring can have potential energy if the spring is stretched or compressed. If the spring is stretched, then there is an attractive energy between the balls, and if it is compressed, then there is a repulsive energy. If you give a spring
relax, then in both cases the potential energy of the balls will turn into kinetic energy.
Chemical energy is one form of potential energy. It exists because the different parts of the molecules attract or repel each other. The more potential energy an object has, the less stable it is. . In reactions, chemical energy can be released in the form of heat energy.
It is practically impossible to determine the absolute content of energy in a molecule. And therefore we are talking only about relative potential energy molecules. It is convenient to represent the relative potential energy of molecules in the form of relative enthalpy. The difference in the relative enthalpies of reactants and products during reactions is denoted by DH°. For exothermic reactions, DH° has negative meaning, and endothermic - positive. When a hydrogen molecule is formed from atoms, heat is released, and when a hydrogen molecule is split into atoms, heat must be supplied:
H + H ¾® H¾H DH° = –104 kcal/mol (–435 kJ/mol)
H-H ¾® H + H DH° = +104 kcal/mol (+435 kJ/mol)
1 kcal = 4.184 kJ
When a chlorine molecule is formed from energy atoms, less energy is released than when a hydrogen molecule is formed:
Сl + Cl ¾® Сl¾Cl DH° = –58 kcal/mol
Cl-Cl ¾® Cl + Cl DH° = +58 kcal/mol
Table 3
Bond energies, kcal/mol.
If we compare the energies of single, double and triple carbon-carbon bonds, we can see that the energy of a double bond is less than two times, and a triple bond is less than three times the energy of a simple C-C connections. Therefore, the transformation of multiple bonds into simple ones, for example, during polymerization, is accompanied by the release of energy.
Bond energy (Е), kcal/mol 88 146 200
For other elements, the opposite pattern is most often observed. For example, when going from simple to double and triple nitrogen-nitrogen bonds, their energy more than doubles and triples.
Bond energy (Е), kcal/mol 38 100 226
Thus, for carbon, the formation of carbon chains is beneficial, and for nitrogen, diatomic molecules. Nitrogen-nitrogen chains can consist of no more than four atoms.
Molecular architecture
from "Molecule Mystery"
Organic chemistry is the chemistry of compounds of carbon. The compounds of carbon and hydrogen are called hydrocarbons. There are thousands of hydrocarbons, many of which are found in natural gas and oil. The simplest hydrocarbon is methane, the main constituent of natural gas. The methane molecule consists of one carbon atom and four hydrogen atoms.Chemists love visuals, so they created structural formulas and various spatial models. Particularly successful are hemispherical (calot - from the French Kalotte - a round cap) molecular models according to Stuart and Brigleb, which take into account the radius of action of individual atoms.
To build a three-dimensional model of the methane molecule, let's take a carbon atom and four hydrogen atoms from the model box and connect them so that the carbon atom is surrounded by four hydrogen atoms. The assembled model has a tetrahedral shape.
As the chain length of alkanes increases, a transition is observed from gases to liquids, and then to waxy bodies. As follows from Table. 1, the first members of the alkane series - from C to C 4 - at normal pressure and room temperature - gases pentane and hexane are light liquids, oily products from C15 to C 7, and solids from C.
Already on the simplest organic compounds a regular relationship between the structure and properties is found. An experienced chemist can often draw a conclusion about its properties and action based on the construction plan of a compound.
Carbon can also form compounds with hydrogen that contain less hydrogen than alkanes. For example, the total composition of ethylene (ethene) is C2H4, i.e. it contains two less hydrogen atoms than ethane. Here, a double bond is formed between both carbon atoms, and both atoms are already in a different state than in alkanes, the bond angle is not tetrahedral (109 ° 28) - it is 120 °.
To build the ethylene molecule, we must take double bonded carbon atoms from the model box (Figure 4).
In acetylene (ethyne C2H2) there are two carbon atoms with a triple bond, the bond angle is 180°. Using carbon atoms in the form of corresponding hemispheres, we construct a model of the acetylene molecule (Fig. 4).
Of great importance are cyclic compounds, such as cycloalkaves (cycloparaffins), such as cyclopentane and cyclohexane, which are representatives of the previously mentioned naphthenes.
The most important among the so-called aromatic cyclic compounds is benzene. Chemists of the last century wondered for a long time how a substance is built that corresponds to the composition of SHF. This compound, obviously unsaturated, behaves very differently from ethylene, propylene, or acetylene. Illumination came to A. Kekule. It is said that he dreamed of a snake biting its own tail. So he imagined the ring structure of the 26th benzene molecule.
The formula in best corresponds to the special interaction (state) of bonds in benzene using a circle in a hexagon, it expresses that three pairs of electrons benzene ring united in a single sextet. For a clearer understanding in Fig. 4 shows a hemispherical model of benzene.
In the case of homonuclear molecules, we carried out the $AO$ combination using the rule according to which the orbitals of the same energy interact most strongly. In heteronuclear molecules of the $AB$ type, the energy levels of the $A$ and $B$ atoms are not the same, so it is difficult to unequivocally state which orbitals will be combined. For the $LiH$ case, this is shown in Fig. one.
Figure 1. Energy levels of $AO$, $Li$ and $H$
Polyatomic chemical particles(molecules, radicals, ions) with the general molecular formula $B_n$, containing one central $A$ atom, two or more terminal $B$ atoms and, as a result, only $A-B \sigma $ bonds.
geometric shape particles $AB_n$ is derived from the method valence bonds, i.e., from the stereochemical arrangement of the axes of the valence hybrid orbitals of the central atom $A$, and hence the $\sigma $-bonds $A-B$.
Hybrid orbitals help to understand the spatial structure of molecules, for example, why the water molecule has an angular configuration, ammonia has a pyramidal, and methane has a tetrahedral configuration.
Consideration of the connection between hybridization and the shape of molecules
Beryllium hydride, $BeH_2$, has a linear structure. To construct its localized bonding molecular orbitals, first, two equivalent valence orbitals of the $Be$ atom are formed, directed towards two hydrogen atoms, $H_a$ and $H_b$, respectively. This is done by hybridization, or mixing (composing a linear combination), $2s-$ and $2p$-orbitals of $Be$, resulting in two equivalent "$sp$-hybrid" orbitals. One of these hybrid orbitals, $sp_a$, is directed towards the $H_a$ atom and strongly overlaps with the $1s_a$ orbital. Another hybrid orbital, $sp_b$, is directed towards the $H_b$ atom and strongly overlaps with the $1s_b$ orbital. With this line of reasoning, two $BeH_2$ bonding molecular orbitals are obtained by constructing two equivalent linear combinations, each of which is localized between two atoms:
These localized molecular orbitals are shown in Fig. 2. Four valence electrons are located on them, forming two localized bonding electron pairs, in accordance with the Lewis bond structure for $BeH_2$. Each of the linear $sp$-hybrid orbitals is half $p$-character and half $s$-character, and two $sp$-orbitals allow the central $Be$ atom in $BeH_2$ to attach two hydrogen atoms to itself.
Let us now consider the $BH_3$ molecule (which is observed in mass spectrometric experiments and is a fragment of the $B_2H_6$ molecule). In this molecule, three hydrogen atoms are attached to the central boron atom. According to the theory of localized molecular orbitals, the bond in this molecule occurs as a result of hybridization of the $2s$ orbital and two $2p$ orbitals of the boron atom to form three equivalent $sp^2$ hybrid orbitals. Each hybrid orbital has one third $s$ character and two thirds $p$ character. Since any two $p$-orbitals lie in the same plane, and the $s$-orbital has no spatial orientation, three $sp^2$-hybrid orbitals lie in the same plane. These three $sp^2$ hybrid orbitals overlap with the three hydrogen $1s$ orbitals to form three equivalent localized bonding orbitals. Each of these bonding $(sp^2+1s)$-orbitals is occupied in the $BH_3$ molecule by a pair of electrons, as shown schematically in Fig. 4. Based on the concept of hybrid orbitals, it can be predicted that the $BH_3$ molecule should have a planar trigonal structure. The angle between the internuclear axes $H-B-H$, called the bond angle $H-B-H$, should be $120^\circ$.
Figure 2. Bonding pairs of electrons on localized bonds of the $BeH_2$ molecule, formed with the participation of equivalent hybrid $sp$-orbitals of the $Be$ atom. Each $sp$-orbital $Be$ forms a localized bonding molecular orbital with the $1s$-orbital of the hydrogen atom
Figure 3. Mutual overlap of orbitals. Hybrid orbitals: a - overlapping $s$-orbitals; b - overlap of $s-$ orbitals; c - overlap of $p-$ orbitals; d - $p$-hybrid orbital; e - $sp^2$-hybrid orbitals; e - $sp^3$-hybrid orbitals
Figure 4. Electron pairs socialized on localized bonds in $BH_3$
Methane, $CH_4$, has four equivalent hydrogen atoms attached to the central carbon atom. To combine with four hydrogen atoms, carbon has to use all of its valence orbitals. By hybridizing one $2s-$ and three $2p$ orbitals, four equivalent $sp^3$ hybrid orbitals can be obtained. Each $sp^3$ hybrid orbital has one quarter $s$ character and three quarters $p$ character. All four $sp^3$-orbitals are directed to the vertices of a regular tetrahedron, therefore $sp^3$-orbitals are sometimes called tetrahedral hybrids. As a result of the overlap of each $sp^3$-hybrid orbital with the $1s-$ orbital of the hydrogen atom, four localized bonding orbitals are formed. The best overlap between the $sp^3$ and $1s$ orbitals is obtained by placing four hydrogen atoms at the vertices of a regular tetrahedron, as shown in Fig. 5 (which shows a cube whose alternating vertices form the vertices of the mentioned tetrahedron). The methane molecule has eight valence electrons (four from the carbon atom and one from each of the four hydrogen atoms), which must be placed in four localized bonding orbitals. These eight electrons form four equivalent localized bonding electron pairs, shown schematically in Fig. 5.
The structure of the $CH_4$ molecule was determined by various experimental methods. All the obtained data lead to the conclusion about the tetrahedral structure of the $CH_4$ molecule (Fig. 6), in full agreement with the predictions of the theory of localized molecular orbitals. Valence angle$H-C-H$ is $109.5^\circ$ and the $C-H$ bond length is $1.093 A$.
Figure 5. Electron pairs socialized on localized bonds in $CH_4$
Figure 6. Tetrahedral molecular structure of $CH_4$
Continuation. For the beginning, see № 15, 16/2004
Lesson 5
atomic orbitals of carbon
A covalent chemical bond is formed using common bonding electron pairs of the type:
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.
For agreement electronic structure carbon atom and the valency of this element use the concept of excitation of the carbon atom. 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:
On 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 p x, 2r y and 2 pz according to the axes along which these orbitals are located.
When chemical bonds are formed, the electron orbitals acquire the same shape. Yes, in saturated hydrocarbons mixed one s-orbital and three R-orbitals of a carbon atom to form four identical (hybrid) sp 3-orbitals:
This is - sp 3 - hybridization.
Hybridization– alignment (mixing) of atomic orbitals ( s and R) with the formation of new atomic orbitals, called hybrid orbitals.
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.
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 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.
Covalent bond formed by lateral overlap R-orbitals of neighboring carbon atoms is called pi()-bond.
Education
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Due to less overlap of orbitals, the -bond is less strong than the -bond.
sp-Hybridization is a 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 placed perpendicular to each other.
directions - connections. On the image sp-orbitals are shown along the axis y, and the unhybridized two
R-orbitals - along the axes X and z.
The triple carbon-carbon bond CC consists of a -bond that occurs when overlapping
sp-hybrid orbitals, and two -bonds.
The relationship between such parameters of the carbon atom as the number of attached groups, the type of hybridization and the types of chemical bonds formed is shown in Table 4.
Table 4
Covalent bonds of carbon
| Number of groups related with carbon |
Type hybridization |
Types participating chemical bonds |
Examples of compound formulas |
|---|---|---|---|
| 4 | sp 3 | Four - connections | |
| 3 | sp 2 | Three - connections and one is connection |
|
| 2 | sp | Two - connections and two connections |
H-CC-H |
Exercises.
1. What electrons of atoms (for example, carbon or nitrogen) are called unpaired?
2. What does the concept of "shared electron pairs" mean in compounds with a covalent bond (for example, CH 4 or H 2 S )?
3. What are the electronic states of atoms (for example, C or N ) are called basic, and which are excited?
4. What do the numbers and letters mean in the electronic formula of an atom (for example, C or N )?
5. What is an atomic orbital? How many orbitals are in the second energy level of a C atom and how do they differ?
6. What is the difference between hybrid orbitals and the original orbitals from which they were formed?
7. What types of hybridization are known for the carbon atom and what are they?
8. draw a picture spatial arrangement orbitals for one of the electronic states of the carbon atom.
9. What chemical bonds are called and what? Specify-and-connections in connections:
10. For the carbon atoms of the compounds below, indicate: a) the type of hybridization; b) types of its chemical bonds; c) bond angles.
Answers to exercises for topic 1
Lesson 5
1. Electrons that are one per orbital are called unpaired electrons. For example, in the electron diffraction formula of an excited carbon atom, there are four unpaired electrons, and the nitrogen atom has three:
2. Two electrons involved in the formation of one chemical bond, called common electron pair. Usually, before the formation of a chemical bond, one of the electrons of this pair belonged to one atom, and the other electron belonged to another atom:
3. The electronic state of the atom, in which the order of filling of electronic orbitals is observed: 1 s 2 , 2s 2 , 2p 2 , 3s 2 , 3p 2 , 4s 2 , 3d 2 , 4p 2 etc. are called main state. AT excited state one of the valence electrons of the atom occupies a free orbital with a higher energy, such a transition is accompanied by the separation of paired electrons. Schematically it is written like this:
Whereas in the ground state there were only two valence unpaired electrons, in the excited state there are four such electrons.
5.
An atomic orbital is a function that describes the density of an electron cloud at each point in space around the nucleus of a given atom. There are four orbitals on the second energy level of the carbon atom - 2 s, 2p x, 2r y,
2pz. These orbitals are:
a) the shape of the electron cloud ( s- ball, R- dumbbell);
b) R-orbitals have different orientations in space - along mutually perpendicular axes x, y and z, they are denoted p x, r y,
pz.
6.
Hybrid orbitals differ from the original (non-hybrid) orbitals in shape and energy. For example, s-orbital - the shape of a sphere, R- symmetrical figure eight, sp-hybrid orbital - asymmetric figure eight.
Energy Differences: E(s) < E(sp) < E(R). Thus, sp-orbital - an orbital averaged in shape and energy, obtained by mixing the initial s-
and p-orbitals.
7. Three types of hybridization are known for the carbon atom: sp 3 , sp 2 and sp (see the text of lesson 5).
9.
-bond - a covalent bond formed by frontal overlapping of orbitals along a line connecting the centers of atoms.
-bond - a covalent bond formed by lateral overlap R-orbitals on either side of the line connecting the centers of atoms.
- Bonds are shown by the second and third lines between the connected atoms.
