As already mentioned, a common electron pair that carries out a covalent bond can be formed due to unpaired electrons present in unexcited interacting atoms. This occurs, for example, during the formation of such molecules as H2, HC1, Cl2. Here each of the atoms has one unpaired electron; when two such atoms interact, a common electron pair is created - a covalent bond arises.
An unexcited nitrogen atom has three unpaired electrons:
Therefore, due to unpaired electrons, the nitrogen atom can participate in the formation of three covalent bonds. This is what happens, for example, in N 2 or NH 3 molecules, in which the covalence of nitrogen is 3.
However, the number of covalent bonds can be more number unpaired electrons present in an unexcited atom. So, in the normal state, the outer electron layer of the carbon atom has a structure that is depicted by the diagram:
Due to the presence of unpaired electrons, a carbon atom can form two covalent bonds. Meanwhile, carbon is characterized by compounds in which each of its atoms is associated with neighboring atoms by four covalent bonds (for example, CO 2 , CH 4, etc.). This is possible due to the fact that, with the expenditure of some energy, one of the 2x electrons present in the atom can be transferred to the sublevel 2 R as a result, the atom goes into an excited state, and the number of unpaired electrons increases. Such an excitation process, accompanied by the "pairing" of electrons, can be represented by the following scheme, in which the excited state is marked with an asterisk next to the symbol of the element:
Now there are four unpaired electrons in the outer electron layer of the carbon atom; therefore, an excited carbon atom can participate in the formation of four covalent bonds. In this case, an increase in the number of created covalent bonds is accompanied by the release more energy than it takes to bring an atom into an excited state.
If the excitation of an atom, leading to an increase in the number of unpaired electrons, is associated with very large energy costs, then these costs are not compensated by the energy of the formation of new bonds; then such a process as a whole turns out to be energetically unfavorable. Thus, oxygen and fluorine atoms do not have free orbitals in the outer electron layer:
Here, an increase in the number of unpaired electrons is possible only by transferring one of the electrons to the next energy level, i.e. into a state 3s. However, such a transition is associated with a very large expenditure of energy, which is not covered by the energy released during the formation of new bonds. Therefore, due to unpaired electrons, an oxygen atom can form no more than two covalent bonds, and a fluorine atom can form only one. Indeed, these elements are characterized by a constant covalence equal to two for oxygen and one for fluorine.
Atoms of elements of the third and subsequent periods have in the outer electron layer "i-sublevel, to which, upon excitation, they can go s- and p-electrons of the outer layer. Therefore, additional possibilities for increasing the number of unpaired electrons appear here. Thus, the chlorine atom, which in the unexcited state has one unpaired electron
can be transferred with the expenditure of some energy into excited states (SI), characterized by three, five or seven unpaired electrons: 
Therefore, unlike the fluorine atom, the chlorine atom can participate in the formation of not only one, but also three, five, or seven covalent bonds. So, in hydrochloric acid HClO 2, the covalence of chlorine is three, in chloric acid HClO 3 - five, and in perchloric acid HClO 4 - seven. Similarly, the sulfur atom, which also has an unoccupied 36Ciod level, can go into excited states with four or six unpaired electrons and, therefore, participate in the formation of not only two, as in oxygen, but also four or six covalent bonds. This can explain the existence of compounds in which sulfur exhibits a covalence equal to four (SO 2 , SCl 4) or six (SF 6).
In many cases, covalent bonds also arise due to the paired electrons present in the outer electron layer of the atom. Consider, for example, the electronic structure of the ammonia molecule:
Here, the dots denote the electrons that originally belonged to the nitrogen atom, and the crosses denote those that belonged to the hydrogen atoms. Of the eight outer electrons of the nitrogen atom, six form three covalent bonds and are common to the nitrogen atom and hydrogen atoms. But two electrons belong only to nitrogen and form lone electron pair. Such a pair of electrons can also participate in the formation of a covalent bond with another atom if there is a free orbital in the outer electron layer of this atom. An unfilled ls-orbital is present, for example, in the hydrogen ion H +, which is generally devoid of electrons:
Therefore, when an NH 3 molecule interacts with a hydrogen ion, a covalent bond arises between them; the lone pair of electrons of the nitrogen atom becomes common to two atoms, resulting in the formation of an ion ammonium NH4:
Here, a covalent bond arose due to a pair of electrons originally belonging to one atom (donor electron pair), and a free orbital of another atom (acceptor electron pair). This way of forming a covalent bond is called donor-acceptor. In the considered example, the electron pair donor is a nitrogen atom, and the acceptor is a hydrogen atom.
Experience has shown that four N-H bonds in the ammonium ion are equivalent in all respects. It follows from this that the bond formed by the donor-acceptor method does not differ in its properties from the covalent bond created due to the unpaired electrons of the interacting atoms.
Another example of a molecule in which there are bonds formed by the donor-acceptor method is the nitric oxide (I) N 2 O molecule.
Earlier structural formula this compound was depicted as follows:
According to this formula, the central nitrogen atom is connected to neighboring atoms by five covalent bonds, so that there are ten electrons (five electron pairs) in its outer electron layer. But such a conclusion contradicts the electronic structure of the nitrogen atom, since its outer L-layer contains only four orbitals (one 5- and three p-orbitals) and cannot contain more than eight electrons. Therefore, the above structural formula cannot be considered correct.
Let us consider the electronic structure of nitric oxide (I), and the electrons of individual atoms will be alternately denoted by dots or crosses. The oxygen atom, which has two unpaired electrons, forms two covalent bonds with the central nitrogen atom:
Due to the unpaired electron remaining at the central nitrogen atom, the latter forms a covalent bond with the second nitrogen atom:
Thus, the outer electron layers of the oxygen atom and the central nitrogen atom are filled: stable eight-electron configurations are formed here. But only six electrons are located in the outer electron layer of the extreme nitrogen atom; this atom can therefore be an acceptor of another electron pair. The central nitrogen atom adjacent to it has an unshared electron pair and can act as a donor. This leads to the formation of another covalent bond between nitrogen atoms by the donor-acceptor method:
Now each of the three atoms that make up the N 2 O molecule has a stable eight-electron structure of the outer layer. If the covalent bond formed by the donor-acceptor method is denoted, as is customary, by an arrow pointing from the donor atom to the acceptor atom, then the structural formula of nitric oxide (I) can be represented as follows:
Thus, in nitric oxide (I), the covalence of the central nitrogen atom is four, and the extreme one is two.
The considered examples show that atoms have various possibilities for the formation of covalent bonds. The latter can be created both at the expense of unpaired electrons of an unexcited atom, and at the expense of unpaired electrons that appear as a result of the excitation of an atom (the “pairing” of electron pairs), and, finally, by the donor-acceptor method. However, the total number of covalent bonds that a given atom can form is limited. It is defined total number valence orbitals, i.e. those orbitals, the use of which for the formation of covalent bonds turns out to be energetically favorable. Quantum-mechanical calculation shows that such orbitals include S- and p-orbitals of the outer electron layer and d-orbitals of the previous layer; in some cases, as we have seen with the examples of chlorine and sulfur atoms, the b/ orbitals of the outer layer can also be used as valence orbitals.
Atoms of all elements of the second period have four orbitals in the outer electron layer in the absence of ^-orbitals in the previous layer. Therefore, the valence orbitals of these atoms can accommodate no more than eight electrons. This means that the maximum covalence of the elements of the second period is four.
Atoms of elements of the third and subsequent periods can be used to form covalent bonds not only s- And R-, but also ^-orbitals. Compounds of ^-elements are known in which s- And R-orbitals of the outer electron layer and all five
The ability of atoms to participate in the formation of a limited number of covalent bonds is called satiety covalent bond.
- A covalent bond formed by the donor-acceptor process is sometimes briefly referred to as a donor-acceptor bond. However, this term should not be understood as a special type of bond, but only a certain way of forming a covalent bond.
There are two main ways (mechanisms) for the formation of a covalent bond.
1) Spinvalent (exchange) mechanism : an electron pair forming a bond is formed due to unpaired electrons present in unexcited atoms.
However, the number of covalent bonds can be greater than the number of unpaired electrons. For example, in the unexcited state (also called the ground state), the carbon atom has two unpaired electrons, but it is characterized by compounds in which it forms four covalent bonds. This becomes possible as a result of the excitation of the atom. In this case, one of the s-electrons goes to the p-sublevel:
An increase in the number of created covalent bonds is accompanied by the release of more energy than is spent on excitation of the atom. Since the valency of an atom depends on the number of unpaired electrons, excitation leads to an increase in valence. In atom nitrogen, oxygen, fluorine, the number of unpaired electrons does not increase, because there are no free orbitals within the second level, and the movement of electrons to the third quantum level requires much more energy than that which would be released during the formation of additional bonds. Thus, when an atom is excited, the transitions of electrons to free orbitals are possible only within one energy level.
Elements of the 3rd period - phosphorus, sulfur, chlorine - can show a valency equal to the group number. This is achieved by excitation of atoms with the transition of 3s and 3p electrons to vacant orbitals of the 3d sublevel:
P* 1s 2 2s 2 2p 6 3s 1 3p 3 3d 1 (valence 5)
S* 1s 2 2s 2 2p 6 3s 1 3p 3 3d 2 (valency 6)
Cl* 1s 2 2s 2 2p 6 3s 1 3p 3 3d 3 (valence 7)
In the above electronic formulas for excited atoms, the sublevels containing only unpaired electrons are underlined. Using the example of the chlorine atom, it is easy to show that the valence can be variable:
Unlike chlorine, the valency of the F atom is constant and equal to 1, because there are no d-sublevel orbitals and other vacant orbitals on the valence (second) energy level.
2) Donor-acceptor mechanism : covalent bonds are formed due to the paired electrons present on the outer electron layer of the atom. In this case, the second atom must have a free orbital on the outer layer. For example, the formation of an ammonium ion from an ammonia molecule and a hydrogen ion can be displayed by the scheme:
An atom that provides its electron pair for the formation of a covalent bond is called a donor, and an atom that provides an empty orbital is called an acceptor. covalent bond formed in this way is called a donor-acceptor bond. In the ammonium cation, this bond is absolutely identical in its properties to the other three covalent bonds formed by the exchange method.
Hybridization of atomic orbitals
To explain the difference between the bond angles in H 2 O (104.5) and NH 3 (107.3) molecules from 90, it should be taken into account that the stable state of the molecule corresponds to its geometric structure with the lowest potential energy. Therefore, during the formation of a molecule, the shape and mutual arrangement of atomic electron clouds change in comparison with their shape and arrangement in free atoms. As a result, a more complete overlap of orbitals is achieved during the formation of a chemical bond. Such deformation of electron clouds requires energy, but more complete overlap leads to the formation of a stronger bond, and in general there is a gain in energy. This explains the emergence of hybrid orbitals.
The shape of the hybrid orbital can be determined mathematically by adding the wave functions of the original orbitals:
As a result of adding the wave functions of s- and p-orbitals, taking into account their signs, it turns out that the density of the electron cloud (value 2) is increased on one side of the nucleus, and lowered on the other.
In general, the process of hybridization includes the following stages: excitation of an atom, hybridization of the orbitals of an excited atom, formation of bonds with other atoms. The energy costs for the first two stages are compensated by the energy gain in the formation of stronger bonds with hybrid orbitals. The type of hybridization is determined by the type and number of orbitals involved in it.
Examples of different types of hybridization of s- and p-orbitals are considered below.
Hybridization of one s- and one p-orbital (sp-hybridization) occurs, for example, in the formation of beryllium hydride, beryllium, zinc, cadmium-mercury halides. The atoms of these elements in the normal state have two paired s-electrons in the outer layer. As a result of excitation, one of the s-electrons passes into the p-state - two unpaired electrons appear, one of which is an s-electron and the other is a p-electron. When a chemical bond is formed, these two different orbitals are converted into two identical hybrid orbitals. The total number of orbitals during hybridization does not change . Two sp-hybrid orbitals are directed at an angle of 180º to each other and form two bonds (Figure 2):
Figure 2 - Overlapping sp-orbitals of beryllium and p-orbitals of chlorine in the BeCl 2 molecule
Experimental determination of the structure of molecules BeГ 2 , ZnГ 2 , CdГ 2 , HgГ 2 (Г-halogen) showed that these molecules are linear, and both metal bonds with halogen atoms have the same length.
Hybridization of one s- and two p-orbitals (sp 2 -hybridization) takes place, for example, in the formation of boron compounds. An excited boron atom has three unpaired electrons - one s-electron and two p-electrons. From three orbitals, three equivalent sp 2 -hybrid orbitals are formed, located in the same plane at an angle of 120 to each other (Figure 3). Indeed, as experimental studies show, the molecules of such boron compounds as BG 3 (G-halogen), B (CH 3) 3 - trimethylboron, B (OH) 3 - boric acid, have a planar structure. In this case, three boron bonds in these molecules have the same length and are located at an angle of 120.
Figure 3– Overlapping sp 2 orbitals of boron and p orbitals of chlorine in the BCl 3 molecule
Hybridization of one s- and three p-orbitals (sp 3 -hybridization) is typical, for example, for carbon and its analogs, silicon and germanium. In this case, four hybrid sp3 orbitals are located at an angle of 10928 to each other; they are directed to the vertices of the tetrahedron (in the molecules CH 4 , CCl 4 , SiH 4 , GeBr 4 and others). The bond angles in the H 2 O (104.5º) and NH 3 (107.3º) molecules do not exactly correspond to the mutual arrangement of “pure” p-orbitals (90º). This is due to some contribution of s-electrons to the formation of a chemical bond. This contribution is nothing but hybridization. Valence electrons in these molecules occupy four orbitals, which are close to sp 3 hybrid. The slight difference between the bond angles and the tetrahedral angles of 109º28" is explained, according to Gillespie's theory, by the fact that the unshared hybrid orbitals occupy a larger volume in space.
In many molecules, the central atom does not undergo hybridization. Thus, the bond angles in the molecules H 2 S, PH 3, etc. are close to 90, i.e. the formation of bonds occurs with the participation of “pure” p-orbitals located at right angles to each other. 
USE OF NEW INFORMATION
TECHNOLOGIES IN CHEMISTRY LESSONS
Time is quickly running forward, and if earlier the school needed to create a theoretical base and educational and methodological support, now there is everything necessary to increase the efficiency of its work. And this is the great merit of the national project "Education". Of course, we, teachers, experience great difficulties in terms of mastering modern technologies. Our inability to work with a computer affects, and it takes a lot of time to master it. But still very interesting and exciting! Moreover, the result is obvious. Children are interested in the lessons, a variety of classes are very fast and informative.
People often think that chemistry is harmful and dangerous. We often hear: “Environmentally friendly products!”, “I heard that you are being poisoned with chemicals!”... But this is not true! We, teachers of chemistry, are faced with the task of convincing schoolchildren that chemistry is a creative science, that it is the productive force of society, and its products are used in all industries, Agriculture and without chemicals it is impossible further development civilization.
The widespread introduction of chemicals, substances, methods and technological methods requires highly educated specialists with a solid base of chemical knowledge. To do this, our school has a specialized chemical and biological class, which provides high-quality training for schoolchildren to continue chemical education. In order for students in high school to choose this particular profile, in the 9th grade there is an elective course "Chemistry in everyday life", the purpose of which is to help children get acquainted with professions directly related to the subjects of chemistry and biology. Even if students do not choose a chemical-biological profile in high school, the knowledge about substances that they constantly encounter in everyday life will be useful in life.
In the classroom of the elective course, the first place is given to lectures. In preparing for them, I use information Internet resources. Many illustrations, diagrams, video collections, laboratory work materials, slides are displayed on the screen, and based on them, I lead my story. My explanation technology has changed significantly. The children are very interested, they listen to the story with great attention and desire.
Chemistry is an experimental science. A large amount of time is devoted to laboratory work. But it happens that some reagents are not in the laboratory, and a virtual laboratory comes to the rescue. With the help of a special program, students can conduct a virtual experiment. Children study the effect of synthetic detergents on different kinds tissues, water solubility of mineral fertilizers, the medium of their solution, the qualitative composition of food (carbohydrates, proteins, fats). With the help of a computer, they keep their own experimental diary, where they fix the topic laboratory work, their observations, conclusions on the correct use of these substances in everyday life. The advantages of a virtual laboratory are safety, no need for laboratory equipment, and time costs are minimal.
At the end of the course, students must pass a test on any topic studied. They are faced with the task of choosing in what form to sum up. The most traditional is a test in the form of an abstract, message or report. For their preparation, children use materials from Internet resources. In this, of course, I help them: I clearly set the task, formulating the questions that the students must answer, and indicate the address of the site with information on the relevant topic.
But this form is already a little outdated, and some guys began to choose project activities. They work individually, in groups, in teams. The search for information is not complete without using the capabilities of the Internet. Before releasing them to the free search, I give them an orientation: search technique, keywords, phrases, names of search engines that may be useful to work with, addresses of sites on the Internet.
Children also choose a test in the form of a game, tasks and exercises for which they develop themselves. It can be a test-turntable, “Smarts and smarts”, “How to become a millionaire?”, “What? Where? When?", various puzzles.
I also arrange a presentation of the resulting product with the involvement of remote technologies. By posting the results of activities on the Internet on the website of a school or class, students get the opportunity to evaluate their work not only with the help of their classmates, but also with the help of children and teachers from other schools, discuss these results, look at them with different eyes.
From the point of view of the new media pedagogy, we live in an extremely interesting time. The rapid introduction of modern technologies forces us to approach old positions in a new way. Pre-profile training at our school exists for four years, and each time I review the course of the lessons, because. new perspectives are opening up, fruitful links are emerging between traditional teaching methods and the new tasks of society, information and knowledge. Indeed, media education has become part of general education. At the same time, the children develop communication skills, interest in new technologies, enthusiasm, individual activity, creativity, they actively cooperate, exchange their own opinions.
I am convinced that the use of information technology can provide a developed learning culture. This is success in teaching and learning. Apply information technology! Move from old forms of classes that have lost their effectiveness to newer, more advanced and modern ones!
The use of new information technologies in educational process can be illustrated by the example of one of the lessons in general chemistry in the 11th grade.
The mechanism of formation and properties of a covalent bond
The purpose of the lesson. Recall from the course of the 8th grade the mechanism of formation of a covalent bond, to study the donor-acceptor mechanism and the properties of a covalent bond.
Equipment. Table of electronegativity chemical elements, codograms of st- and l-bonds, training disk " general chemistry» from a series of training programs of Cyril and Methodius with diagrams and models of molecules, ball-and-stick models of molecules, a work card with tasks and tests, an interactive whiteboard, a computer, tasks for consolidating and controlling knowledge with remote control.
During the classes
The lecture is conducted with the help of the training disk "General Chemistry".
Repetition of the material covered
Recall with students, due to which a bond is formed between the atoms of non-metals. Complete tasks 1, 2 on the work card (see appendix).
Learning new material
Covalent bond formation mechanism:
a) exchange (for example, H 2, Cl 2, HC1);
b) donor-acceptor (for example, NH 4 C1).
Immediately, students write down their homework in the margins: Depict the formation of the hydronium ion H 3 ABOUT + from H ion + and water molecules.
Types of covalent bonds: polar and non-polar (according to the composition of the molecule).
Properties of a covalent bond.
multiplicity(single, one and a half, double, triple).
Bond energy is the amount of energy released during the formation chemical bond or spent on breaking it.
Link length is the distance between the nuclei of atoms in a molecule.
The energy and length of the bond are interconnected. Show by example how these properties are interconnected, how they affect the strength of the molecule (project onto the board):
With an increase in the number of bonds between atoms in a molecule, the bond length decreases, and its energy increases, for example (project onto a board):Saturability- this is the ability of atoms to form a certain and limited number of bonds. Show with ball-and-stick examples
molecules Cl 2, H 2 O, CH 4, HNO 3.
Orientation. Consider patterns of overlapping electron clouds during the formation of σ- and π-bonds, project onto the board (Fig.).
Fix tasks 6, 7 on the work card (see appendix).
Small break!
1. Let's start the list in order,
Because the first element.
(It forms, by the way, water -
very important point).
Let's imagine a molecule
Convenient formula H 2 .
Let's add -
There is no lighter substance in the world!
2. N 2 is a nitrogen molecule.
Known to be colorless
gas. A lot of knowledge, but let's
Let's replenish them anyway.
3. He is everywhere and everywhere:
And in stone, in air, in water,
He is in the morning dew
And blue in the sky.
(Oxygen.)
4. Mushroom pickers found a small swamp in the forest, from which gas bubbles escaped in places. The match ignited the gas, and a faint flame began to wander through the swamp. What is this gas? (Methane)
Continuation of the lesson.
Polarizability is the ability of a covalent bond to change its polarity under the influence of an external electric field(pay attention to such different concepts as bond polarity and molecular polarizability).
Consolidation of the studied material
Control on the studied topic is carried out using remote controls.
The survey is conducted within 3 minutes, 10 questions at the price of one point, 30 seconds are given for the answer, the questions are projected onto an interactive whiteboard. When scoring 9-10 points - score "5", 7-8 points - score "4", 5-6 points - score "3".
Questions for consolidation
1. The bond that is formed due to common electron pairs is called:
a) ionic; b) covalent; c) metal.
2. A covalent bond is formed between atoms:
a) metals; b) non-metals; c) metal and non-metal.
3. The mechanism for the formation of a covalent bond due to the lone electron pair of one atom and the free orbital of another is called:
a) donor-acceptor; b) inert; c) catalytic.
4. Which of the molecules has a covalent bond?
a) Zn; b) Cu O; c) NH3.
5. The multiplicity of bonds in a nitrogen molecule is equal to:
a) three; b) two; c) unit.
6. The smallest bond length in a molecule:
a) H2S; b) SF6; c) SO 2 ; d) SOr
7. When electron clouds overlap along the axis connecting the nuclei of interacting atoms, the following is formed:
a) σ-bond; b) π bond; c) ρ-bond.
8. A nitrogen atom has a possible number of unpaired electrons:
a) 1; b) 2; at 3.
9. Bond strength increases in the series:
a) H 2 O - H 2 S; 6) NH 3 - PH 3; c) CS 2 - C O 2; d) N 2 - O 2
10. The hybrid s-orbital has the form:
a) a ball b) wrong eight; c) the correct eight.
The results are immediately displayed on the screen, we make a report on each question.
Analysis of homework (see appendix - work card), § 6 of the textbook by O.S. Gabrielyan, G. Glysov “Chemistry. Grade 11 ”(M .: Drofa, 2006), abstract in a notebook.
Application
work card
1. Match the names of the substance and the type of connection.
1) Potassium chloride;
2) oxygen;
3) magnesium;
4) carbon tetrachloride.
a) Covalent non-polar;
b) ionic;
c) metal;
d) covalent polar.
2. Between the atoms of which elements will the chemical bond have an ionic character?
a) NnO; b) Si and C1; c) Na and O; d) P and Br.
3. The bond length is expressed in:
a) nm; b) kg; c) j; d) m 3.
4. Where is the chemical bond the strongest: in the Cl 2 or O 2 molecule?
5. In which molecule is the strength of the hydrogen bond greater: H 2 O or H 2 S?
6. Continue the sentence: “The bond formed by the overlapping of electron clouds along the line connecting the nuclei of atoms is called ............................................ ......",
7. Sketch the patterns of overlapping electron orbitals during the formation of a π bond.
8. Homework. "General chemistry in tests, tasks, exercises" O.S. Gabrielyan (Moscow: Drofa, 2003), work 8A, option 1, 2.
The idea of the formation of a chemical bond with the help of a pair of electrons belonging to both connecting atoms was put forward in 1916 by the American physical chemist J. Lewis.
A covalent bond exists between atoms both in molecules and in crystals. It occurs both between identical atoms (for example, in H 2, Cl 2, O 2 molecules, in a diamond crystal), and between different atoms (for example, in H 2 O and NH 3 molecules, in SiC crystals). Almost all bonds in molecules organic compounds are covalent (C-C, C-H, C-N, etc.).
There are two mechanisms for the formation of a covalent bond:
1) exchange;
2) donor-acceptor.
Exchange mechanism for the formation of a covalent bondis that each of the connecting atoms provides for the formation of a common electron pair (bond) by one unpaired electron. The electrons of the interacting atoms must have opposite spins.
Consider, for example, the formation of a covalent bond in a hydrogen molecule. When hydrogen atoms approach each other, their electron clouds penetrate each other, which is called the overlap of electron clouds (Fig. 3.2), the electron density between the nuclei increases. The nuclei are attracted to each other. As a result, the energy of the system decreases. With a very strong approach of atoms, the repulsion of nuclei increases. Therefore, there is an optimal distance between the nuclei (bond length l) at which the system has a minimum energy. In this state, energy is released, called the binding energy E St.
Rice. 3.2. Scheme of overlapping electron clouds during the formation of a hydrogen molecule
Schematically, the formation of a hydrogen molecule from atoms can be represented as follows (a dot means an electron, a bar means a pair of electrons):
H + H→H: H or H + H→H - H.
IN general view for AB molecules of other substances:
A + B = A: B.
Donor-acceptor mechanism of covalent bond formationconsists in the fact that one particle - the donor - presents an electron pair for the formation of a bond, and the second - the acceptor - a free orbital:
A: + B = A: B.
donor acceptor
Consider the mechanisms of formation of chemical bonds in the ammonia molecule and the ammonium ion.
1. Education
The nitrogen atom has two paired and three unpaired electrons in its outer energy level:
The hydrogen atom on the s - sublevel has one unpaired electron.
In the ammonia molecule, the unpaired 2p electrons of the nitrogen atom form three electron pairs with the electrons of 3 hydrogen atoms:
In the NH 3 molecule, 3 covalent bonds are formed by the exchange mechanism.
2. The formation of a complex ion - an ammonium ion.
NH 3 + HCl = NH 4 Cl or NH 3 + H + = NH 4 +
The nitrogen atom has a lone pair of electrons, i.e. two electrons with antiparallel spins in the same atomic orbital. The atomic orbital of the hydrogen ion does not contain electrons (a vacant orbital). When an ammonia molecule and a hydrogen ion approach each other, the lone pair of electrons of the nitrogen atom and the vacant orbital of the hydrogen ion interact. The unshared pair of electrons becomes common for nitrogen and hydrogen atoms, a chemical bond arises according to the donor-acceptor mechanism. The nitrogen atom of the ammonia molecule is the donor, and the hydrogen ion is the acceptor:
It should be noted that in the NH 4 + ion all four bonds are equivalent and indistinguishable, therefore, in the ion the charge is delocalized (dispersed) over the entire complex.
The considered examples show that the ability of an atom to form covalent bonds is determined not only by one-electron, but also by 2-electron clouds or by the presence of free orbitals.
According to the donor-acceptor mechanism, bonds are formed in complex compounds: - ; 2+ ; 2- etc.
A covalent bond has the following properties:
- satiety;
- orientation;
- polarity and polarizability.
Ticket number 11
Ticket number 12
Ticket number 13
Ticket number 14
Ticket number 15.
EXAMINATION TICKET No. 11
Redox reactions. The oxidation state of an element. Examples of oxidizing and reducing agents.
The method of valence bonds (MVS). Exchange and donor-acceptor mechanisms of covalent bond formation.
Answer:
Redox reactions(OVR) - reactions that go with a change in s.d. atoms. Redox reactions are chemical reactions that occur with a change in the oxidation states of the atoms that make up the reactants, realized by the redistribution of electrons between the oxidizing atom and the reducing atom.
Oxidation state(s.d.) - the charge that is attributed to the atom, considering it an ion
Oxidizer (Ox) accepts electrons.
Restorer (Red) - donates electrons
Ox 1 + Red 2 Red 1 + Ox 2
Ox1 + ne– → Red1
Cu2+ + 2e– → Cu0
CuSO 4 + Zn → ZnSO 4 + Cu
Red2–ne– → Ox2
Zn0 – 2e– → Zn2+
Valence bond method
1927 - Heitler and London Quantum-mechanical calculation of the hydrogen molecule
Communication formation mechanisms
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Exchange A + BA : IN |
Donor-acceptor A : +VA : IN BF 3 + F – – : NH 3 + H + + |
The mechanism of formation of a covalent bond.
MVS makes it possible to distinguish three mechanisms for the formation of a covalent bond: exchange, donor-acceptor, and dative.
exchange mechanism. It includes those cases of the formation of a chemical bond, when each of the two bonded atoms allocates one electron for socialization, as if exchanging them. To bind the nuclei of two atoms, the electrons must be in the space between the nuclei. This area in the molecule is called the binding area (the area where the electron pair is most likely to stay in the molecule). In order for the exchange of unpaired electrons between atoms, the overlap of atomic orbitals is necessary. This is the action of the exchange mechanism for the formation of a covalent chemical bond. Atomic orbitals can only overlap if they have the same symmetry properties about the internuclear axis.
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Donor-acceptor and dative mechanisms.
The donor-acceptor mechanism is associated with the transfer of a lone pair of electrons from one atom to a vacant one. atomic orbital another atom. For example, the formation of an ion -:
The vacant p-AO in the boron atom in the BF 3 molecule accepts a pair of electrons from the fluoride ion (donor). In the resulting anion, four B-F covalent bonds are equivalent in length and energy. In the original molecule, all three B-F bonds were formed by the exchange mechanism.
Atoms, the outer shell of which consists only of s- or p-electrons, can be either donors or acceptors of the lone pair of electrons. Atoms that have valence electrons on the d-AO can simultaneously act as both donors and acceptors. To distinguish between these two mechanisms, the concepts of the dative mechanism of bond formation were introduced.
Examination ticket number 12
The second law of thermodynamics. Entropy, her physical meaning and calculation methods. The change in the entropy of the system as a probabilistic criterion for the direction of the process.
Osmosis. osmotic pressure. Van't Hoff's law for solutions of non-electrolytes.
Answer:
Second law of thermodynamics
IN isolated system, a spontaneous process is possible only with an increase in entropy.
