Electronic graphic circuit o2. Electronic formulas and graphic diagrams of the structure of electronic layers of atoms. Orbits: technical data

Let's look at how an atom is built. Keep in mind that we will talk exclusively about models. In practice, atoms are much more complex structure. But thanks to modern developments, we are able to explain and even successfully predict properties (even if not all). So what is the structure of an atom? What is it “made” of?

Planetary model of the atom

It was first proposed by the Danish physicist N. Bohr in 1913. This is the first theory of atomic structure based on scientific facts. In addition, it laid the foundation for modern thematic terminology. In it, electrons-particles produce rotational movements around the atom according to the same principle as the planets around the Sun. Bohr suggested that they could exist exclusively in orbits located at a strictly defined distance from the nucleus. The scientist could not explain why this was so, from a scientific standpoint, but such a model was confirmed by many experiments. Integer numbers were used to designate orbits, starting with one, which was numbered closest to the nucleus. All these orbits are also called levels. The hydrogen atom has only one level, on which one electron rotates. But complex atoms also have levels. They are divided into components that combine electrons with similar energy potential. So, the second already has two sublevels - 2s and 2p. The third already has three - 3s, 3p and 3d. And so on. First, the sublevels closer to the core are “populated,” and then the distant ones. Each of them can only hold a certain number of electrons. But this is not the end. Each sublevel is divided into orbitals. Let's make a comparison with ordinary life. The electron cloud of an atom is comparable to a city. Levels are streets. Sublevel - a private house or an apartment. Orbital - room. Each of them “lives” one or two electrons. They all have specific addresses. This was the first diagram of the structure of the atom. And finally, about the addresses of electrons: they are determined by sets of numbers that are called “quantum”.

Wave model of the atom

But over time, the planetary model was revised. A second theory of atomic structure was proposed. It is more advanced and allows you to explain the results of practical experiments. The first one was replaced by the wave model of the atom, which was proposed by E. Schrödinger. Then it was already established that an electron can manifest itself not only as a particle, but also as a wave. What did Schrödinger do? He applied an equation that describes the motion of a wave in Thus, one can find not the trajectory of an electron in an atom, but the probability of its detection at a certain point. What unites both theories is that elementary particles are located at specific levels, sublevels and orbitals. This is where the similarity between the models ends. Let me give you one example: in wave theory, an orbital is a region where an electron can be found with a 95% probability. The rest of the space accounts for 5%. But in the end it turned out that the structural features of atoms are depicted using the wave model, despite the fact that the terminology used is common.

The concept of probability in this case

Why was this term used? Heisenberg formulated the uncertainty principle in 1927, which is now used to describe the movement of microparticles. It is based on their fundamental difference from ordinary physical bodies. What is it? Classical mechanics assumed that a person could observe phenomena without influencing them (observation of celestial bodies). Based on the data obtained, it is possible to calculate where the object will be at a certain point in time. But in the microcosm things are necessarily different. So, for example, it is now not possible to observe an electron without influencing it due to the fact that the energies of the instrument and the particle are incomparable. This leads to changes in its location of the elementary particle, state, direction, speed of movement and other parameters. And it makes no sense to talk about exact characteristics. The uncertainty principle itself tells us that it is impossible to calculate the exact trajectory of an electron around the nucleus. You can only indicate the probability of finding a particle in a certain area of ​​space. This is the peculiarity of the structure of atoms of chemical elements. But this should be taken into account exclusively by scientists in practical experiments.

Atomic composition

But let's concentrate on the entire subject matter. So, in addition to the well-considered electron shell, the second component of the atom is the nucleus. It consists of positively charged protons and neutral neutrons. We are all familiar with the periodic table. The number of each element corresponds to the number of protons it contains. The number of neutrons is equal to the difference between the mass of an atom and its number of protons. There may be deviations from this rule. Then they say that an isotope of the element is present. The structure of an atom is such that it is “surrounded” by an electron shell. usually equals the number of protons. The mass of the latter is approximately 1840 times greater than that of the former, and is approximately equal to the weight of the neutron. The radius of the nucleus is about 1/200,000 the diameter of the atom. It itself has a spherical shape. This, in general, is the structure of the atoms of chemical elements. Despite the difference in mass and properties, they look approximately the same.

Orbits

When talking about what an atomic structure diagram is, one cannot remain silent about them. So, there are these types:

1. s. They have a spherical shape.
2. p. They look like three-dimensional figure eights or a spindle.
3. d and f. They have a complex shape that is difficult to describe in formal language.

An electron of each type can be found with a 95% probability in the corresponding orbital. The information presented must be treated calmly, since it is rather an abstract mathematical model than a physical reality. But with all this, it has good predictive power regarding the chemical properties of atoms and even molecules. The further a level is located from the nucleus, the more electrons can be placed on it. Thus, the number of orbitals can be calculated using a special formula: x 2. Here x is equal to the number of levels. And since up to two electrons can be placed in an orbital, ultimately the formula for their numerical search will look like this: 2x 2.

Orbits: technical data

If we talk about the structure of the fluorine atom, it will have three orbitals. They will all be filled. The energy of orbitals within one sublevel is the same. To designate them, add the layer number: 2s, 4p, 6d. Let's return to the conversation about the structure of the fluorine atom. It will have two s- and one p-sublevel. It has nine protons and the same number of electrons. First one s-level. That's two electrons. Then the second s-level. Two more electrons. And 5 fills the p-level. This is his structure. After reading the next subheading, you can do it yourself necessary actions and make sure of it. If we talk about which fluorine also belongs, it should be noted that they, although in the same group, are completely different in their characteristics. Thus, their boiling point ranges from -188 to 309 degrees Celsius. So why were they united? All thanks chemical properties. All halogens, and fluorine to the greatest extent, have the highest oxidizing ability. They react with metals and can spontaneously ignite at room temperature without any problems.

How are orbits filled?

By what rules and principles are electrons arranged? We suggest that you familiarize yourself with the three main ones, the wording of which has been simplified for better understanding:

1. Principle of least energy. Electrons tend to fill orbitals in order of increasing energy.
2. Pauli's principle. One orbital cannot contain more than two electrons.
3. Hund's rule. Within one sublevel, electrons first fill empty orbitals, and only then form pairs.

The structure of the atom will help in filling it out and in this case it will become more understandable in terms of image. Therefore, when practical work When constructing circuit diagrams, you need to keep it at hand.

Example

In order to summarize everything that has been said within the framework of the article, you can draw up a sample of how the electrons of an atom are distributed among their levels, sublevels and orbitals (that is, what the configuration of levels is). It can be depicted as a formula, an energy diagram, or a layer diagram. There are very good illustrations here, which, upon careful examination, help to understand the structure of the atom. So, the first level is filled in first. It has only one sublevel, in which there is only one orbital. All levels are filled sequentially, starting with the smallest. First, within one sublevel, one electron is placed in each orbital. Then pairs are created. And if there are free ones, a switch to another filling subject occurs. And now you can find out for yourself what the structure of the nitrogen or fluorine atom is (which was considered earlier). It may be a little difficult at first, but you can use the pictures to guide you. For clarity, let's look at the structure of the nitrogen atom. It has 7 protons (together with neutrons that make up the nucleus) and the same number of electrons (which make up the electron shell). The first s-level is filled in first. It has 2 electrons. Then comes the second s-level. It also has 2 electrons. And the other three are placed on the p-level, where each of them occupies one orbital.

Conclusion

As you can see, the structure of the atom is not such a difficult topic (if you approach it from the perspective of a school chemistry course, of course). And understanding this topic is not difficult. Finally, I would like to tell you about some features. For example, speaking about the structure of the oxygen atom, we know that it has eight protons and 8-10 neutrons. And since everything in nature tends to balance, two oxygen atoms form a molecule, where two unpaired electrons form a covalent bond. Another stable oxygen molecule, ozone (O3), is formed in a similar way. Knowing the structure of the oxygen atom, you can correctly compose formulas oxidative reactions, which involve the most common substance on Earth.

To correctly depict the electronic configurations of atoms, you need to answer the questions: 1. How to determine the total number of electrons in an atom? 2. What is the maximum number of electrons at levels and sublevels? 3. What is the order of filling sublevels and orbitals? 3

Electronic configurations (using the example of a hydrogen atom) 1. Scheme electronic structure The diagram of the electronic structure of atoms shows the distribution of electrons across energy levels 2. Electronic formula 1s 1, where s is the designation of the sublevel; 1 - number of electrons Electronic formulas atoms show the distribution of electrons among energy sublevels 3. Electron graphic formula Electron graphic formulas of atoms show the distribution of electrons among orbitals and electron spins 4

2. Based on the sample, compose the electronic formula of aluminum. The order of filling the energy levels in the atom. 1s 2, 2s 2, 2p 6, 3s 2, 3p 1 6 Aluminum has 13 electrons. The first sublevel in an atom to be filled is the 1s sublevel. It can have a maximum of 2 electrons, mark them and subtract them from the total number of electrons. There are 11 electrons left to place. The next 2s sublevel is filled; it can have 2 electrons. There are 9 electrons left to place. The next 2p sublevel is filled; it can have 6 electrons. Next, we fill the 3s sublevel. We reached the 3p sublevel, there can be a maximum of 6 electrons on it, but there is only 1 left, so we place it. 1s = Al s2s2s 2p2p 3p - 2 = - 6 = - 2 = 9 3 1

3. Determine: Are the energy levels in order? If the levels are in order, then leave them that way. If the levels are not in order, then rewrite them, arranging them in ascending order. No. The 4s and 3d sublevels are out of order. We need to rewrite and arrange them in ascending order. 7 Cr 24 1s 2 2p62p6 3s 2 4s 2 3p 6 3d 4 2s22s2 1s 2 2p62p6 3s 2 4s 2 3p 6 3d 4 2s22s2

Rules for drawing up an electron graphic diagram Each sublevel has a certain number of orbitals. Each orbital can contain no more than two electrons. If there are two electrons in an orbital, then they must have different spins (arrows point in different sides). 8 s p d f Let's start drawing up an electronic graphic diagram
5. Geographical journey Determine in which groups of the periodic table the chemical elements are located, the electronic formulas of the atoms of which are given in the first column of the table. The letters corresponding to the correct answers will give the name of the country. 10 JAMAICA Electronic formulas Groups IIIIIIIVVVIVII 1s 2 2s 1 JAGLRKAO 1s 2 2s 2 2p 6 3s 2 3p 5 VISNPDM 1s 2 2s 2 2p 6 3s 2 3p 4 EFTZJAO 1s 2 2s 2 2p 4 GRISJK 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 KUERMIP 1s 2 2s 2 2p 6 3s 1 ANDLOZHL

2. Structure of nuclei and electron shells of atoms

2.7. Distribution of electrons in an atom

The state of electrons in an atom is indicated using a specific form of notation. For example, for the helium atom we have:

The distribution of electrons in an atom is indicated using:

A) electronic circuits, in which only the number of electrons in each layer is noted. For example: Mg 2e, 8e, 2e; Cl 2e, 8e, 7e.

Graphical electronic circuits are often used, for example, for the chlorine atom:

b) electronic configurations; in this case, the number of the layer (level), the nature of the sublevels and the number of electrons on them are shown. For example:
Mg 1s 2 2s 2 2p 6 3s 2 ;

V) electronic graphic circuits, in which orbitals are depicted, for example, in the form of a cage, and electrons are represented by arrows (Fig. 2.6).

Rice. 2.6. Electron graphic diagram for a magnesium atom

In addition to full formulas for electronic configurations, abbreviated ones are widely used. In this case, the portion of the electron configuration corresponding to the noble gas is indicated by the noble gas symbol in square brackets. For example: 12 Mg3s 2, 19 K4s 1.

There are certain principles and rules for filling energy levels and sublevels with electrons:

1. The principle of the minimum total energy of an atom, according to which the population of the JSC with electrons occurs in such a way that the total energy of the atom is minimal. The following sequence of filling the AO was experimentally established:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p ... .

2. There can be no more than two electrons on one AO, and their spins in this case must be antiparallel.

3. Within a given energy sublevel, electrons fill the AO gradually, first one at a time (first all vacant, and then two at a time), and the orientation of all unpaired electrons must be the same, i.e. such

but not like that

In almost any atom, only s- and p-AOs are external (Fig. 2.7), therefore the outer electron layer of any atom cannot contain more than eight electrons. The outer electron layer, containing eight electrons (in the case of helium, two) is called complete.

Rice. 2.7. Electron graphic diagrams for atoms K (a) and S (b)

The energy values ​​of different energy sublevels for different atoms are not constant, but depend on the charge of the nucleus Z of the element’s atom: for atoms of elements with Z = 1–20 E 3 d > E 4 s and E 3 d > E 4 p; for atoms of elements with Z ≥ 21 vice versa: E 3 d< E 4 s и Е 3 d < E 4 p (рис. 2.8). Кроме того, чем больше Z , тем меньше различаются подуровни по энергии, а кривые, выражающие зависимость энергии подуровней от Z , пересекаются.

Rice. 2.8. Diagram of energy sublevels of atoms of elements with Z = 1–20 (a), Z ≥ 21 (b)

The electronic configurations of the atoms (ground state) of K and Ca are as follows (see Fig. 2.8):

19 K: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 ,

20 Ca: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 .

Starting from scandium (Z = 21), the 3d sublevel is filled, and 4s electrons remain in the outer layer. The general electronic formula of atoms of elements from Sc to Zn is 3d 1−10 4s 1−2. For example:

21 Sc: 3d 1 4s 2,

25 Mn: 3d 5 4s 2,

28 Ni: 3d 8 4s 2 .

30 Zn: 3d 10 4s 2 .

For chromium and copper, a breakthrough (dip) of the 4s electron to the 3d sublevel is observed: Cr - 3d 5 4s 1, Cu - 3d 10 4s 1. Such a jump from the ns - to the (n − 1)d sublevel is also observed in atoms of other elements (Mo, Ag, Au, Pt) and is explained by the proximity of the energies of the ns - and (n − 1)d sublevels, as well as the stability of half and completely filled d-sublevels.

Further in the 4th period, after 10 d-elements, p-elements follow from Ga ( 3d 10 4s 2 4p 1) to Kr ( 3d 10 4s 2 4p 6).

The formation of d-element cations is associated with the loss of first external ns -, then (n − 1)d -electrons, for example:

Ti: 3d 2 4s 2 → − 2 e − Ti 2+ : 3d 2 → − 1 e − Ti 3+ : 3d 1

Mn: 3d 5 4s 2 → − 2 e − Mn 2+ : 3d 5 → − 2 e − Mn 4+ : 3d 3

Note that in formulas for electronic configurations it is customary to first write down all electrons with a lower value of n, and then move on to indicating electrons with a higher value of the principal quantum number. Therefore, the order of filling and the order of recording energy sublevels for 3d elements do not coincide. For example, in the electronic formula of the scandium atom, the 3d orbital is indicated before the 4s orbital, although the 4s orbital is filled first.

A logical question arises: why is the 4s sublevel filled earlier in atoms of 3d elements, although its energy is greater than the energy of the 3d sublevel? Why, for example, does the Sc atom not have the electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 3d 3 in its ground state?

This happens because the ratio of energies of various electronic states of an atom does not always coincide with the ratio of energies of individual energy sublevels. The energy of the 4s sublevel for 3d elements is greater than the energy of the 3d sublevel, but the energy of the state
3d 1 4s 2 is less than the energy of the 3d 3 state.

This is explained by the fact that the interelectron repulsion, and, accordingly, the energy of the entire state for the configuration...3d 3 (with three electrons on the same energy sublevel) is greater than for the configuration...3d 1 4s 2 (with three electrons, located at different energy levels).

Algorithm for composing the electronic formula of an element:

1. Determine the number of electrons in an atom using the Periodic Table of Chemical Elements D.I. Mendeleev.

2. Using the number of the period in which the element is located, determine the number of energy levels; the number of electrons in the last electronic level corresponds to the group number.

3. Divide the levels into sublevels and orbitals and fill them with electrons in accordance with the rules for filling orbitals:

It must be remembered that the first level contains a maximum of 2 electrons 1s 2, on the second - a maximum of 8 (two s and six R: 2s 2 2p 6), on the third - a maximum of 18 (two s, six p, and ten d: 3s 2 3p 6 3d 10).

• Principal quantum number n should be minimal.
• First to fill s- sublevel, then р-, d- b f- sublevels.
• Electrons fill the orbitals in order of increasing energy of the orbitals (Klechkovsky's rule).
• Within a sublevel, electrons first occupy free orbitals one by one, and only after that they form pairs (Hund’s rule).
• There cannot be more than two electrons in one orbital (Pauli principle).

Examples.

1. Let's create an electronic formula for nitrogen. Nitrogen is number 7 on the periodic table.

2. Let's create the electronic formula for argon. Argon is number 18 on the periodic table.

1s 2 2s 2 2p 6 3s 2 3p 6.

3. Let's create the electronic formula of chromium. Chromium is number 24 on the periodic table.

1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5

Energy diagram of zinc.

4. Let's create the electronic formula of zinc. Zinc is number 30 on the periodic table.

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10

Please note that part of the electronic formula, namely 1s 2 2s 2 2p 6 3s 2 3p 6, is the electronic formula of argon.

The electronic formula of zinc can be represented as:

Electronic configurations of atoms

The total number of electrons in an atom is determined by the charge of its nucleus, i.e., the proton number. It is equal to the atomic number of the element. Electrons, depending on their energy, are distributed in the atom into energy levels and sublevels, each of which consists of a certain number of orbitals.

The distribution of electrons is expressed using the electron formulas (or electron configurations) of the atom. For example, hydrogen, an element with atomic number 1, has an electronic formula: 1H 1s1. In this formula, the number of the energy level is written with a number, followed by a letter indicating the type of sublevel, and finally, the number at the top right indicates the number of electrons in that sublevel.

Schematically, the electronic structure of an atom is depicted using an electron graphic diagram, in which orbitals are represented as cells and electrons as arrows.

The electron graphic diagram of the hydrogen atom is written as follows:

To correctly depict electronic formulas, you must follow several basic rules.

1st rule: The distribution of electrons in an atom in the ground (most stable) state is determined by the principle of minimum energy: the ground state of the atom corresponds to the lowest possible energy levels and sublevels.

Therefore, electrons (in atoms of elements of the first three periods) fill the orbitals in order of increasing their energy:

1s→2s→2p→3s→3p

2nd rule: Each orbital can contain a maximum of two electrons, with opposite spins.

Thus, helium 2He, next to hydrogen, has the electronic formula:

2Not 1s2,

Since the first electron layer can only contain two electrons, this layer in the helium atom is complete and therefore very stable.

For atoms of elements of the second period, the second energy level is filled, which can contain no more than 8 electrons. First, electrons fill the 2s orbital (for lithium and beryllium atoms):

Since the 2s orbital is filled, the fifth electron of the boron atom B occupies one of the three 2p orbitals. Electronic formula of the boron atom:

and the electronic graphic diagram:

Note that the 2p sublevel is depicted close to the 2s sublevel, but slightly higher. This emphasizes his belonging to the same level (second) and at the same time a greater supply of energy.

3rd rule. Sets the order in which orbitals of one sublevel are filled. Electrons of one sublevel first fill the orbitals one at a time (i.e., all empty), and if the number of electrons is greater than the number of orbitals, then two at a time. Therefore, the electronic formulas of carbon and nitrogen atoms are:

6C 1s22s22p2 and 7N 1s22s22p3

and electronic graphic circuits:

For oxygen, fluorine and neon atoms, the number of electrons increases, and they are forced to be placed in two p-orbitals of the second energy level:

6O 1s22s22p4; 6F 1s22s22p5; 6Ne 1s22s22p6

Electron graphic diagrams of the atoms of these elements:

The electronic configuration of the outer layer 2s22p6 corresponds to its complete filling and is therefore stable.

The third electron layer begins to form in the atoms of elements of the third period. First, the s-sublevel of sodium and magnesium is filled with electrons:

11Na 1s22s22p63s1 12Mg 1s22s22p63s2

and then the p-sublevel for aluminum, silicon, chlorine and argon:

18Ar 1s22s22p63s23p6

Electron graphic diagram for the argon atom:

An argon atom has 8 electrons in its outer electron layer. Consequently, it is complete, since in an atom of any element at the outer energy level there can be a maximum of no more than 8 electrons.

The construction of the third electronic layer does not end there. In accordance with the formula 2n2, it can contain 18 electrons: 8 in the s- and p-sublevels and 10 in the d-sublevel. This sublevel will be formed among the elements of the fourth period. But first, the first two elements of the fourth period - potassium and calcium - have a fourth electronic layer, which opens with the s-sublevel (the energy of the 4s sublevel is slightly less than that of the 3d sublevel:

19K 1s22s22p63s23p64s1 and 19Са 1s22s22p63s23p64s2

Only after this will the d-sublevel of the third, now pre-external, energy level begin to be filled with electrons. Electronic configuration of the scandium atom:

21Sc 1s22s22p63s23p64s23d1,

titanium atom:

21Ti 1s22s22p63s23p64s23d2,

etc., up to zinc. The electronic configuration of its atom is:

21Zn 1s22s22p63s23p64s23d10,

and the electronic graphic diagram:

Since only the orbitals of the third and fourth energy levels of elements of the fourth period are filled with electrons, completely filled levels (in this case, the first and second) are usually not indicated on electron graphic diagrams. Instead, in electronic formulas the symbol of the nearest element VIII A-group with completely filled energy s- and p-sublevels is written: for example, the electronic formula of chlorine is 3s23p5, zinc is 3d104s2, and antimony is 51Sb -4d105s25p3

In addition to electronic formulas and electronic graphic diagrams, electronic diagrams of atoms are sometimes used, in which only the number of electrons at each energy level (electronic layer) is indicated:

The electronic structure of an atom is determined by the charge of its nucleus, which is equal to the atomic number of the element in the periodic table.

The distribution of electrons across energy levels, sublevels and orbitals is displayed using electronic formulas and electron graphic diagrams, as well as electronic diagrams of atoms.

The outer electron layer of an atom of any element can contain no more than 8 electrons. 3.2. Types of Chemical Bonds

Covalent bond– the most general type of chemical bond that arises due to the socialization of an electron pair through exchange mechanism, when each of the interacting atoms supplies one electron, or donor-acceptor mechanism, if an electron pair is transferred for common use by one atom (donor) to another atom (acceptor) (Fig. 3.2).

A classic example of a nonpolar covalent bond (the electronegativity difference is zero) is observed in homonuclear molecules: H–H, F–F. The energy of a two-electron two-center bond lies in the range of 200–2000 kJ∙mol –1.

When a heteroatomic covalent bond is formed, an electron pair is shifted to a more electronegative atom, which makes the bond polar. The ionicity of a polar bond as a percentage is calculated by the empirical relation 16(χ A – χ B) + 3.5(χ A – χ B) 2, where χ A and χ B are the electronegativity of atoms A and B of the AB molecule. Except polarizability covalent bond has the property saturation– the ability of an atom to form as many covalent bonds as it has energetically accessible atomic orbitals. About the third property of a covalent bond - focus- will be discussed below (see. valence bond method).

Ionic bond– a special case of covalent, when the resulting electron pair completely belongs to a more electronegative atom, which becomes an anion. The basis for identifying this bond as a separate type is the fact that compounds with such a bond can be described in an electrostatic approximation, considering the ionic bond to be due to the attraction of positive and negative ions. The interaction of ions of the opposite sign does not depend on the direction, and Coulomb forces do not have the property of saturation. Therefore, each ion in an ionic compound attracts such a number of ions of the opposite sign that a crystal lattice of an ionic type is formed. There are no molecules in an ionic crystal. Each ion is surrounded by a certain number of ions of a different sign (the coordination number of the ion). Ion pairs can exist in the gaseous state as polar molecules. In the gaseous state, NaCl has a dipole moment of ~3∙10 –29 C∙m, which corresponds to a displacement of 0.8 electron charge per bond length of 0.236 nm from Na to Cl, i.e. Na 0.8+ Cl 0.8– .

The metallic bond arises as a result of partial delocalization of valence electrons, which move quite freely in the metal lattice, electrostatically interacting with positively charged ions. The binding forces are not localized or directed, and delocalized electrons cause high thermal and electrical conductivity.

Hydrogen bond. Its formation is due to the fact that, as a result of a strong displacement of an electron pair towards an electronegative atom, a hydrogen atom, which has an effective positive charge, can interact with another electronegative atom (F, O, N, less often Cl, Br, S). The energy of such electrostatic interaction is 20–100 kJ∙mol –1. Hydrogen bonds can be intra- and intermolecular. An intramolecular hydrogen bond is formed, for example, in acetylacetone and is accompanied by ring closure (Fig. 3.3).

Molecules carboxylic acids in non-polar solvents they dimerize due to two intermolecular hydrogen bonds (Fig. 3.4).

The hydrogen bond plays an extremely important role in biological macromolecules, such inorganic compounds as H 2 O, H 2 F 2, NH 3. Due to hydrogen bonds, water is characterized by such high melting and boiling temperatures compared to H 2 E (E = S, Se, Te). If there were no hydrogen bonds, then water would melt at –100 °C and boil at –80 °C.

Van der Waals (intermolecular) bonding– the most universal type of intermolecular bond, due to dispersion forces(induced dipole – induced dipole), induction interaction (permanent dipole – induced dipole) and orientational interaction (permanent dipole – permanent dipole). The energy of the van der Waals bond is less than the hydrogen bond and amounts to 2–20 kJ∙mol –1.

Chemical bonding in solids. The properties of solids are determined by the nature of the particles occupying the sites of the crystal lattice and the type of interaction between them.

Solid argon and methane form atomic and molecular crystals, respectively. Since the forces between atoms and molecules in these lattices are of the weak van der Waals type, such substances melt at fairly low temperatures. Most substances that are in liquid and gaseous states at room temperature form molecular crystals at low temperatures.

The melting points of ionic crystals are higher than those of atomic and molecular crystals because the electrostatic forces acting between ions far exceed the weak van der Waals forces. Ionic compounds are harder and more brittle. Such crystals are formed by elements with widely different electronegativities (for example, alkali metal halides). Ionic crystals containing polyatomic ions have lower melting points; so for NaCl t pl. = 801 °C, and for NaNO 3 t pl = 306.5 °C.

In covalent crystals, the lattice is built from atoms connected by a covalent bond, so these crystals have high hardness, melting point and low thermal and electrical conductivity.

Crystal lattices formed by metals are called metallic. The sites of such lattices contain positive metal ions, and the interstices contain valence electrons (electron gas).

Among the metals, d-elements have the highest melting point, which is explained by the presence in the crystals of these elements of a covalent bond formed by unpaired d-electrons, in addition to the metallic bond formed by s-electrons.

Valence bond method(MVS) otherwise called the theory of localized electron pairs, since the method is based on the assumption that the chemical bond between two atoms is carried out using one or more electron pairs that are localized primarily between them. Unlike MMO, in which the simplest chemical bond can be either two- or multi-center, in MBC it is always two-electron and necessarily two-center. The number of elementary chemical bonds that an atom or ion can form is equal to its valency. Just as in MMO, valence electrons take part in the formation of a chemical bond. The wave function that describes the state of the electrons forming a bond is called a localized orbital (LO).

Note that electrons described by LO, in accordance with the principle Pauli must have oppositely directed spins, that is, in the MBC all spins are paired, and all molecules must be diamagnetic. Consequently, MMS fundamentally cannot explain the magnetic properties of molecules.

However, the principle of localized connections has a number of important advantages, one of which is its extreme visibility. MBC predicts quite well, for example, the valence capabilities of atoms and the geometry of the resulting molecule. The last circumstance is associated with the so-called hybridization of AO. It was introduced to explain the fact that two-electron two-center chemical bonds formed by AOs in different energy states have the same energy. Thus, Be*(2s 1 1p 1), B*(2s 1 2p 2), C*(2s 1 2p 3) form two, three and four bonds, respectively, due to the s- and p-orbitals, and therefore one of them should be stronger than others. However, experience shows that in BeH 2, BCl 3, CH 4 all bonds are equivalent. For BeH 2 the bond angle is 180°, for BCl 3 it is 120°, and for CH 4 it is 109°28".

According to the concept of hybridization, chemical bonds are formed by mixed - hybrid orbitals (HO), which are a linear combination of AO of a given atom (s- and p-AO Be, B, C), have the same energy and shape, a certain orientation in space (symmetry ). Thus, the s- and p-orbitals give rise to two sp-GOs located at an angle of 180° relative to each other.

In the CH 4 molecule, hybrid orbitals of four carbon AOs (one s and three p) are called sp 3 orbitals; they are completely equivalent in energy and spatially directed to the vertices of the tetrahedron.

Thus, when one atom forms several bonds, and its valence electrons belong to different orbitals (s and p; s, p and d), to explain the geometry of molecules in the MBC it is necessary to invoke the theory of hybridization of atomic orbitals. The main provisions of the theory are as follows:

The introduction of hybrid orbitals serves to describe directional localized bonds. Hybrid orbitals provide maximum overlap of AOs in the direction of localized σ bonds.

The number of hybrid orbitals is equal to the number of AOs participating in hybridization.

Valence AOs that are close in energy hybridize, regardless of whether they are completely filled, half filled, or empty in the atom.

AOs that have common symmetry characteristics participate in hybridization.

According to table. 3.3 hybrid orbitals give molecules with angles of 180°, 120°, 109° 28", 90°. These are regular geometric figures. Such molecules are formed when all the peripheral atoms in a multi-electron molecule (or ion) are the same and their number coincides with the number of hybrid orbitals However, if the number of hybrid orbitals is greater than the number of bonded atoms, then some of the hybrid orbitals are occupied by electron pairs that are not involved in bond formation - non-binding or unshared electron pairs.

H–Be–H, HC≡CH

H 2 C=CH 2, C 6 H 6, BCl 3

CH 4, CCl 4, H 3 C–CH 3

d 2 sp 3 or sp 3 d 2

As an example, consider the molecules NH 3 and H 2 O. Nitrogen and oxygen atoms are prone to sp 3 hybridization. In nitrogen on sp 3 -GO, in addition to three bonding pairs of electrons forming a bond with three hydrogen atoms, there remains one non-bonding pair. It is this that, occupying one sp 3 -GO, distorts the H–N–H bond angle to 107.3°. In the H 2 O molecule there are two such non-bonding pairs, and the H–O–H angle is 104.5° (Fig. 3.17).

The electrons of bonding and nonbonding pairs interact with each other differently. The stronger the interelectron repulsion, the larger the conventional surface area on the sphere occupied by the electron pair. For a qualitative explanation of experimental facts, it is usually believed that non-bonding pairs occupy a larger volume than bonding ones, and the volume of bonding pairs is smaller, the greater the electronegativity of peripheral atoms (method Gillespie).

Physical properties of metals.

Density. This is one of the most important characteristics of metals and alloys. According to their density, metals are divided into the following groups:

lungs(density not more than 5 g/cm 3) - magnesium, aluminum, titanium, etc.:

heavy- (density from 5 to 10 g/cm 3) - iron, nickel, copper, zinc, tin, etc. (this is the most extensive group);

very heavy(density more than 10 g/cm3) - molybdenum, tungsten, gold, lead, etc.

Table 2 shows the density values ​​of metals. (This and the following tables characterize the properties of those metals that form the basis of alloys for artistic casting).

Table 2. Metal density.

Melting temperature. Depending on the melting point, the metal is divided into the following groups:

fusible(melting point does not exceed 600 o C) - zinc, tin, lead, bismuth, etc.;

medium-melting(from 600 o C to 1600 o C) - these include almost half of the metals, including magnesium, aluminum, iron, nickel, copper, gold;

refractory(more than 1600 o C) - tungsten, molybdenum, titanium, chromium, etc.

Mercury is a liquid.

When making artistic castings, the melting point of the metal or alloy determines the choice of melting unit and refractory molding material. When additives are introduced into a metal, the melting point, as a rule, decreases.

Table 3. Melting and boiling points of metals.

Specific heat. This is the amount of energy required to raise the temperature of a unit mass by one degree. Specific heat capacity decreases with increasing atomic number of an element in the periodic table. The dependence of the specific heat capacity of an element in the solid state on atomic mass is described approximately by the Dulong and Petit law:

m a c m = 6.

Where, m a- atomic mass; c m- specific heat capacity (J/kg * o C).

Table 4 shows the specific heat capacity of some metals.

Table 4. Specific heat capacity of metals.

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