The location of electrons on energy shells or levels is recorded using electronic formulas of chemical elements. Electronic formulas or configurations help to represent the structure of an element's atom.

The structure of the atom

The atoms of all elements consist of a positively charged nucleus and negatively charged electrons that are located around the nucleus.

The electrons are at different energy levels. The farther an electron is from the nucleus, the more energy it has. The size of the energy level is determined by the size of the atomic orbit or orbital cloud. This is the space in which the electron moves.

Rice. 1. The general structure of the atom.

Orbitals can have different geometric configurations:

• s-orbitals- spherical;
• p-, d and f-orbitals- dumbbell-shaped, lying in different planes.

At the first energy level of any atom, there is always an s-orbital with two electrons (an exception is hydrogen). Starting from the second level, the s- and p-orbitals are at the same level.

Rice. 2. s-, p-, d and f-orbitals.

Orbitals exist regardless of the location of electrons on them and can be filled or vacant.

Formula entry

Electronic configurations of atoms of chemical elements are written according to the following principles:

• each energy level corresponds to a serial number, denoted by an Arabic numeral;
• the number is followed by a letter denoting the orbital;
• a superscript is written above the letter, corresponding to the number of electrons in the orbital.

Recording examples:

• calcium -

  1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 ;

• oxygen -

  1s 2 2s 2 2p 4 ;

• carbon-

  1s 2 2s 2 2p 2 .

The periodic table helps to write down the electronic formula. The number of energy levels corresponds to the number of the period. The number of the element indicates the charge of an atom and the number of electrons. The group number indicates how many valence electrons are in the outer level.

Let's take Na as an example. Sodium is in the first group, in the third period, at number 11. This means that the sodium atom has a positively charged nucleus (contains 11 protons), around which 11 electrons are located at three energy levels. There is one electron in the outer level.

Recall that the first energy level contains an s-orbital with two electrons, and the second contains s- and p-orbitals. It remains to fill the levels and get the full record:

11 Na) 2) 8) 1 or 1s 2 2s 2 2p 6 3s 1 .

For convenience, special tables of electronic formulas of the element have been created. In the long periodic table, the formulas are also indicated in each cell of the element.

Rice. 3. Table of electronic formulas.

For brevity, elements are written in square brackets, the electronic formula of which coincides with the beginning of the element formula. For example, the electronic formula of magnesium is 3s 2, neon is 1s 2 2s 2 2p 6. Therefore, the full formula for magnesium is 1s 2 2s 2 2p 6 3s 2 . 4.6. Total ratings received: 195.

The conditional image of the distribution of electrons in the electron cloud by levels, sublevels and orbitals is called the electronic formula of the atom.

Rules based on|based on| which | which | make up | hand over | electronic formulas

1. Principle of minimum energy: the less energy the system has, the more stable it is.

2. Klechkovsky's rule: the distribution of electrons over the levels and sublevels of the electron cloud occurs in ascending order of the sum of the principal and orbital quantum numbers (n + 1). In the case of equality of values ​​(n + 1), the sublevel that has the smaller value of n is filled first.

1 s 2 s p 3 s p d 4 s p d f 5 s p d f 6 s p d f 7 s p d f Level number n 1 2 2 3 3 3 4 4 4 4 5 5 5 5 6 6 6 6 7 7 7 7 Orbital 1* 0 0 1 0 1 2 0 1 2 3 0 1 2 3 0 1 2 3 0 1 2 3 quantum number

n+1| 1 2 3 3 4 5 4 5 6 7 5 6 7 8 6 7 8 9 7 8 9 10

Klechkovsky series

1* - see table No. 2.

3. Hund's rule: when the orbitals of one sublevel are filled, the lowest energy level corresponds to the placement of electrons with parallel spins.

Drafting|Submitting| electronic formulas

Potential row: 1 s 2 s p 3 s p d 4 s p d f 5 s p d f 6 s p d f 7 s p d f

(n+1|) 1 2 3 3 4 5 4 5 6 7 5 6 7 8 6 7 8 9 7 8 9 10

Klechkovsky series

Filling order Electroni 1s 2 2s 2 p 6 3s 2 p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 ..

(n+l|) 1 2 3 3 4 4 5 5 5 6 6 6 7 7 7 7 8.

Electronic formula

(n+1|) 1 2 3 3 4 5 4 5 6 7 5 6 7 8 6 7 8 9 7 8 9 10

Informativeness of electronic formulas

1. The position of the element in the periodic|periodic| system.

2. Possible degrees| element oxidation.

3. The chemical nature of the element.

4. Composition|warehouse| and connection properties of the element.

The position of the element in the periodic|Periodic|D.I. Mendeleev’s system:

a) period number, in which the element is located, corresponds to the number of levels on which the electrons are located;

b) group number, to which this element belongs, is equal to the sum of valence electrons. Valence electrons for atoms of s- and p-elements are electrons of the outer level; for d-elements, these are the electrons of the outer level and the unfilled sublevel of the previous level.

in) electronic family is determined by the symbol of the sublevel to which the last electron enters (s-, p-, d-, f-).

G) subgroup is determined by belonging to the electronic family: s - and p - elements occupy the main subgroups, and d - elements - secondary, f - elements occupy separate sections in the lower part of the periodic system (actinides and lanthanides).

2. Possible degrees| element oxidation.

Oxidation state is the charge that an atom acquires when it gives or gains electrons.

Atoms that donate electrons acquire a positive charge, which is equal to the number of electrons donated (electron charge (-1)

Z E 0 – ne  Z E + n

The atom that donated electrons becomes cation(positive charged ion). The process of removing an electron from an atom is called ionization process. The energy needed to carry out this process is called ionization energy ( Eion, eB).

The first to separate from the atom are electrons of the outer level, which do not have a pair in the orbital - unpaired. In the presence of free orbitals within the same level, under the action of external energy, the electrons that formed pairs at this level are unpaired, and then separated all together. The process of depairing, which occurs as a result of the absorption of a portion of energy by one of the electrons of the pair and its transition to the highest sublevel, is called arousal process.

The largest number of electrons that an atom can donate is equal to the number of valence electrons and corresponds to the number of the group in which the element is located. The charge that an atom acquires after losing all its valence electrons is called the highest degree of oxidation atom.

After release|dismissal| valence level external becomes|becomes| level which|what| preceded valence. This is a level completely filled with electrons, and therefore | and therefore | energy resistant.

Atoms of elements that have from 4 to 7 electrons at the external level achieve an energetically stable state not only by giving up electrons, but also by adding them. As a result, a level (.ns 2 p 6) is formed - a stable inert gas state.

An atom that has attached electrons acquires negativedegreeoxidation- a negative charge, which is equal to the number of received electrons.

Z E 0 + ne  Z E - n

The number of electrons that an atom can attach is equal to the number (8 –N|), where N is the number of the group in which|what| the element is located (or the number of valence electrons).

The process of attaching electrons to an atom is accompanied by the release of energy, which is called c affinity to the electron (Esrodship,eV).

The composition of the atom.

An atom is made up of atomic nucleus and electron shell.

The nucleus of an atom is made up of protons ( p+) and neutrons ( n 0). Most hydrogen atoms have a single proton nucleus.

Number of protons N(p+) is equal to the nuclear charge ( Z) and the ordinal number of the element in the natural series of elements (and in the periodic system of elements).

N(p +) = Z

The sum of the number of neutrons N(n 0), denoted simply by the letter N, and the number of protons Z called mass number and is marked with the letter BUT.

A = Z + N

The electron shell of an atom consists of electrons moving around the nucleus ( e -).

Number of electrons N(e-) in the electron shell of a neutral atom is equal to the number of protons Z at its core.

The mass of a proton is approximately equal to the mass of a neutron and 1840 times the mass of an electron, so the mass of an atom is practically equal to the mass of the nucleus.

The shape of an atom is spherical. The radius of the nucleus is about 100,000 times smaller than the radius of the atom.

Chemical element- type of atoms (set of atoms) with the same nuclear charge (with the same number of protons in the nucleus).

Isotope- a set of atoms of one element with the same number of neutrons in the nucleus (or a type of atoms with the same number of protons and the same number of neutrons in the nucleus).

Different isotopes differ from each other in the number of neutrons in the nuclei of their atoms.

Designation of a single atom or isotope: (E - element symbol), for example: .

The structure of the electron shell of the atom

atomic orbital is the state of an electron in an atom. Orbital symbol - . Each orbital corresponds to an electron cloud.

The orbitals of real atoms in the ground (unexcited) state are of four types: s, p, d and f.

electronic cloud- the part of space in which an electron can be found with a probability of 90 (or more) percent.

Note: sometimes the concepts of "atomic orbital" and "electron cloud" are not distinguished, calling both of them "atomic orbital".

The electron shell of an atom is layered. Electronic layer formed by electron clouds of the same size. Orbitals of one layer form electronic ("energy") level, their energies are the same for the hydrogen atom, but different for other atoms.

Orbitals of the same level are grouped into electronic (energy) sublevels:
s- sublevel (consists of one s-orbitals), symbol - .
p sublevel (consists of three p
d sublevel (consists of five d-orbitals), symbol - .
f sublevel (consists of seven f-orbitals), symbol - .

The energies of the orbitals of the same sublevel are the same.

When designating sublevels, the number of the layer (electronic level) is added to the sublevel symbol, for example: 2 s, 3p, 5d means s- sublevel of the second level, p- sublevel of the third level, d- sublevel of the fifth level.

The total number of sublevels in one level is equal to the level number n. The total number of orbitals in one level is n 2. Accordingly, the total number of clouds in one layer is also n 2 .

Designations: - free orbital (without electrons), - orbital with an unpaired electron, - orbital with an electron pair (with two electrons).

The order in which electrons fill the orbitals of an atom is determined by three laws of nature (formulations are given in a simplified way):

1. The principle of least energy - electrons fill the orbitals in order of increasing energy of the orbitals.

2. Pauli's principle - there cannot be more than two electrons in one orbital.

3. Hund's rule - within the sublevel, electrons first fill free orbitals (one at a time), and only after that they form electron pairs.

The total number of electrons in the electronic level (or in the electronic layer) is 2 n 2 .

The distribution of sublevels by energy is expressed next (in order of increasing energy):

1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p ...

Visually, this sequence is expressed by the energy diagram:

The distribution of electrons of an atom by levels, sublevels and orbitals (the electronic configuration of an atom) can be depicted as an electronic formula, an energy diagram, or, more simply, as a diagram of electronic layers ("electronic diagram").

Examples of the electronic structure of atoms:

Valence electrons- electrons of an atom that can take part in the formation of chemical bonds. For any atom, these are all the outer electrons plus those pre-outer electrons whose energy is greater than that of the outer ones. For example: Ca atom has 4 outer electrons s 2, they are also valence; the Fe atom has external electrons - 4 s 2 but he has 3 d 6, hence the iron atom has 8 valence electrons. The valence electronic formula of the calcium atom is 4 s 2, and iron atoms - 4 s 2 3d 6 .

Periodic system of chemical elements of D. I. Mendeleev
(natural system of chemical elements)

Periodic law of chemical elements(modern formulation): the properties of chemical elements, as well as simple and complex substances formed by them, are in a periodic dependence on the value of the charge from atomic nuclei.

Periodic system- graphical expression of the periodic law.

Natural range of chemical elements- a number of chemical elements, arranged according to the increase in the number of protons in the nuclei of their atoms, or, what is the same, according to the increase in the charges of the nuclei of these atoms. The serial number of an element in this series is equal to the number of protons in the nucleus of any atom of this element.

The table of chemical elements is constructed by "cutting" the natural series of chemical elements into periods(horizontal rows of the table) and groupings (vertical columns of the table) of elements with a similar electronic structure of atoms.

Depending on how elements are combined into groups, a table can be long period(elements with the same number and type of valence electrons are collected in groups) and short-term(elements with the same number of valence electrons are collected in groups).

The groups of the short period table are divided into subgroups ( main and side effects), coinciding with the groups of the long-period table.

All atoms of elements of the same period have the same number of electron layers, equal to the number of the period.

The number of elements in the periods: 2, 8, 8, 18, 18, 32, 32. Most of the elements of the eighth period were obtained artificially, the last elements of this period have not yet been synthesized. All periods except the first start with an alkali metal forming element (Li, Na, K, etc.) and end with a noble gas forming element (He, Ne, Ar, Kr, etc.).

In the short period table - eight groups, each of which is divided into two subgroups (main and secondary), in the long period table - sixteen groups, which are numbered in Roman numerals with the letters A or B, for example: IA, IIIB, VIA, VIIB. Group IA of the long period table corresponds to the main subgroup of the first group of the short period table; group VIIB - secondary subgroup of the seventh group: the rest - similarly.

The characteristics of chemical elements naturally change in groups and periods.

In periods (with increasing serial number)

• the nuclear charge increases
• the number of outer electrons increases,
• the radius of the atoms decreases,
• the bond strength of electrons with the nucleus increases (ionization energy),
• electronegativity increases.
• the oxidizing properties of simple substances are enhanced ("non-metallicity"),
• the reducing properties of simple substances ("metallicity") weaken,
• weakens the basic character of hydroxides and the corresponding oxides,
• the acidic character of hydroxides and corresponding oxides increases.

In groups (with increasing serial number)

• the nuclear charge increases
• the radius of atoms increases (only in A-groups),
• the strength of the bond between electrons and the nucleus decreases (ionization energy; only in A-groups),
• electronegativity decreases (only in A-groups),
• weaken the oxidizing properties of simple substances ("non-metallicity"; only in A-groups),
• the reducing properties of simple substances are enhanced ("metallicity"; only in A-groups),
• the basic character of hydroxides and the corresponding oxides increases (only in A-groups),
• the acidic nature of hydroxides and the corresponding oxides weakens (only in A-groups),
• the stability of hydrogen compounds decreases (their reducing activity increases; only in A-groups).

Tasks and tests on the topic "Topic 9. "The structure of the atom. Periodic law and periodic system of chemical elements of D. I. Mendeleev (PSCE)"."

• Periodic Law - Periodic law and structure of atoms Grade 8–9
  You should know: the laws of filling orbitals with electrons (principle of least energy, Pauli's principle, Hund's rule), the structure of the periodic system of elements.

  You should be able to: determine the composition of an atom by the position of an element in the periodic system, and, conversely, find an element in the periodic system, knowing its composition; depict the structure diagram, the electronic configuration of an atom, ion, and, conversely, determine the position of a chemical element in the PSCE from the diagram and electronic configuration; characterize the element and the substances it forms according to its position in the PSCE; determine changes in the radius of atoms, the properties of chemical elements and the substances they form within one period and one main subgroup of the periodic system.

  Example 1 Determine the number of orbitals in the third electronic level. What are these orbitals?
  To determine the number of orbitals, we use the formula N orbitals = n 2 , where n- level number. N orbitals = 3 2 = 9. One 3 s-, three 3 p- and five 3 d-orbitals.

  Example 2 Determine the atom of which element has the electronic formula 1 s 2 2s 2 2p 6 3s 2 3p 1 .
  In order to determine which element it is, you need to find out its serial number, which is equal to the total number of electrons in the atom. In this case: 2 + 2 + 6 + 2 + 1 = 13. This is aluminum.

  After making sure that everything you need is learned, proceed to the tasks. We wish you success.

  
  Recommended literature:
  • O. S. Gabrielyan and others. Chemistry, 11th grade. M., Bustard, 2002;
  • G. E. Rudzitis, F. G. Feldman. Chemistry 11 cells. M., Education, 2001.

It is written in the form of so-called electronic formulas. In electronic formulas, the letters s, p, d, f denote the energy sublevels of electrons; the numbers in front of the letters indicate the energy level in which the given electron is located, and the index at the top right is the number of electrons in this sublevel. To compose the electronic formula of an atom of any element, it is enough to know the number of this element in the periodic system and fulfill the basic provisions that govern the distribution of electrons in an atom.

The structure of the electron shell of an atom can also be depicted in the form of an arrangement of electrons in energy cells.

For iron atoms, such a scheme has the following form:

This diagram clearly shows the implementation of Hund's rule. At the 3d sublevel, the maximum number of cells (four) is filled with unpaired electrons. The image of the structure of the electron shell in the atom in the form of electronic formulas and in the form of diagrams does not clearly reflect the wave properties of the electron.

The wording of the periodic law as amended YES. Mendeleev : the properties of simple bodies, as well as the forms and properties of the compounds of elements, are in a periodic dependence on the magnitude of the atomic weights of the elements.

Modern formulation of the Periodic Law: the properties of the elements, as well as the forms and properties of their compounds, are in a periodic dependence on the magnitude of the charge of the nucleus of their atoms.

Thus, the positive charge of the nucleus (rather than atomic mass) turned out to be a more accurate argument on which the properties of elements and their compounds depend.

Valence- is the number of chemical bonds that one atom is bonded to another.
The valence possibilities of an atom are determined by the number of unpaired electrons and the presence of free atomic orbitals at the outer level. The structure of the outer energy levels of atoms of chemical elements determines mainly the properties of their atoms. Therefore, these levels are called valence levels. The electrons of these levels, and sometimes of the pre-external levels, can take part in the formation of chemical bonds. Such electrons are also called valence electrons.

Stoichiometric valence chemical element - is the number of equivalents that a given atom can attach to itself, or is the number of equivalents in the atom.

Equivalents are determined by the number of attached or substituted hydrogen atoms, therefore, the stoichiometric valence is equal to the number of hydrogen atoms with which this atom interacts. But not all elements interact freely, but almost everything interacts with oxygen, so the stoichiometric valency can be defined as twice the number of attached oxygen atoms.

For example, the stoichiometric valency of sulfur in hydrogen sulfide H 2 S is 2, in oxide SO 2 - 4, in oxide SO 3 -6.

When determining the stoichiometric valence of an element according to the formula of a binary compound, one should be guided by the rule: the total valency of all atoms of one element must be equal to the total valence of all atoms of another element.

Oxidation state also characterizes the composition of the substance and is equal to the stoichiometric valence with a plus sign (for a metal or a more electropositive element in a molecule) or minus.

1. In simple substances, the oxidation state of elements is zero.

2. The oxidation state of fluorine in all compounds is -1. The remaining halogens (chlorine, bromine, iodine) with metals, hydrogen and other more electropositive elements also have an oxidation state of -1, but in compounds with more electronegative elements they have positive oxidation states.

3. Oxygen in compounds has an oxidation state of -2; the exceptions are hydrogen peroxide H 2 O 2 and its derivatives (Na 2 O 2, BaO 2, etc., in which oxygen has an oxidation state of -1, as well as oxygen fluoride OF 2, in which the oxidation state of oxygen is +2.

4. Alkaline elements (Li, Na, K, etc.) and elements of the main subgroup of the second group of the Periodic system (Be, Mg, Ca, etc.) always have an oxidation state equal to the group number, that is, +1 and +2, respectively .

5. All elements of the third group, except for thallium, have a constant oxidation state equal to the group number, i.e. +3.

6. The highest oxidation state of an element is equal to the group number of the Periodic system, and the lowest is the difference: group number is 8. For example, the highest oxidation state of nitrogen (it is located in the fifth group) is +5 (in nitric acid and its salts), and the lowest is -3 (in ammonia and ammonium salts).

7. The oxidation states of the elements in the compound compensate each other so that their sum for all atoms in a molecule or a neutral formula unit is zero, and for an ion - its charge.

These rules can be used to determine the unknown oxidation state of an element in a compound, if the oxidation states of the others are known, and to formulate multi-element compounds.

Degree of oxidation (oxidation number,) — auxiliary conditional value for recording the processes of oxidation, reduction and redox reactions.

concept oxidation state often used in inorganic chemistry instead of the concept valence. The oxidation state of an atom is equal to the numerical value of the electric charge attributed to the atom, assuming that the electron pairs that carry out the bond are completely biased towards more electronegative atoms (that is, based on the assumption that the compound consists only of ions).

The oxidation state corresponds to the number of electrons that must be added to a positive ion to reduce it to a neutral atom, or taken from a negative ion to oxidize it to a neutral atom:

Al 3+ + 3e − → Al
S 2− → S + 2e − (S 2− − 2e − → S)

The properties of the elements, depending on the structure of the electron shell of the atom, change according to the periods and groups of the periodic system. Since in a number of analogous elements the electronic structures are only similar, but not identical, when moving from one element in a group to another, not a simple repetition of properties is observed for them, but their more or less clearly expressed regular change.

The chemical nature of an element is determined by the ability of its atom to lose or gain electrons. This ability is quantified by the values ​​of ionization energies and electron affinity.

Ionization energy (Ei) is the minimum amount of energy required for the detachment and complete removal of an electron from an atom in the gas phase at T = 0

K without transferring kinetic energy to the released electron with the transformation of the atom into a positively charged ion: E + Ei = E + + e-. The ionization energy is a positive value and has the lowest values ​​for alkali metal atoms and the highest for noble (inert) gas atoms.

Electron affinity (Ee) is the energy released or absorbed when an electron is attached to an atom in the gas phase at T = 0

K with the transformation of the atom into a negatively charged ion without transferring kinetic energy to the particle:

E + e- = E- + Ee.

Halogens, especially fluorine, have the maximum electron affinity (Ee = -328 kJ/mol).

The values ​​of Ei and Ee are expressed in kilojoules per mol (kJ/mol) or in electron volts per atom (eV).

The ability of a bound atom to displace the electrons of chemical bonds towards itself, increasing the electron density around itself is called electronegativity.

This concept was introduced into science by L. Pauling. Electronegativitydenoted by the symbol ÷ and characterizes the tendency of a given atom to attach electrons when it forms a chemical bond.

According to R. Maliken, the electronegativity of an atom is estimated by half the sum of the ionization energies and the electron affinity of free atoms h = (Ee + Ei)/2

In periods, there is a general tendency for an increase in the ionization energy and electronegativity with an increase in the charge of the atomic nucleus; in groups, these values ​​decrease with an increase in the ordinal number of the element.

It should be emphasized that an element cannot be assigned a constant value of electronegativity, since it depends on many factors, in particular, on the valence state of the element, the type of compound in which it is included, the number and type of neighboring atoms.

Atomic and ionic radii. The dimensions of atoms and ions are determined by the dimensions of the electron shell. According to quantum mechanical concepts, the electron shell does not have strictly defined boundaries. Therefore, for the radius of a free atom or ion, we can take theoretically calculated distance from the core to the position of the main maximum density of the outer electron clouds. This distance is called the orbital radius. In practice, the values ​​of the radii of atoms and ions in compounds, calculated from experimental data, are usually used. In this case, covalent and metallic radii of atoms are distinguished.

The dependence of atomic and ionic radii on the charge of the nucleus of an atom of an element and is periodic. In periods, as the atomic number increases, the radii tend to decrease. The greatest decrease is typical for elements of small periods, since the outer electronic level is filled in them. In large periods in the families of d- and f-elements, this change is less sharp, since the filling of electrons in them occurs in the preexternal layer. In subgroups, the radii of atoms and ions of the same type generally increase.

The periodic system of elements is a clear example of the manifestation of various kinds of periodicity in the properties of elements, which is observed horizontally (in a period from left to right), vertically (in a group, for example, from top to bottom), diagonally, i.e. some property of the atom increases or decreases, but the periodicity is preserved.

In the period from left to right (→), the oxidizing and non-metallic properties of the elements increase, while the reducing and metallic properties decrease. So, of all the elements of period 3, sodium will be the most active metal and the strongest reducing agent, and chlorine will be the strongest oxidizing agent.

chemical bond- this is the interconnection of atoms in a molecule, or crystal lattice, as a result of the action of electric forces of attraction between atoms.

This is the interaction of all electrons and all nuclei, leading to the formation of a stable, polyatomic system (radical, molecular ion, molecule, crystal).

Chemical bonding is carried out by valence electrons. According to modern concepts, the chemical bond has an electronic nature, but it is carried out in different ways. Therefore, there are three main types of chemical bonds: covalent, ionic, metallic. Between molecules arises hydrogen bond, and happen van der Waals interactions.

The main characteristics of a chemical bond are:

- bond length - is the internuclear distance between chemically bonded atoms.

It depends on the nature of the interacting atoms and on the multiplicity of the bond. With an increase in the multiplicity, the bond length decreases, and, consequently, its strength increases;

- bond multiplicity - is determined by the number of electron pairs linking two atoms. As the multiplicity increases, the binding energy increases;

- connection angle- the angle between imaginary straight lines passing through the nuclei of two chemically interconnected neighboring atoms;

Binding energy E CB - this is the energy that is released during the formation of this bond and is spent on breaking it, kJ / mol.

covalent bond - A chemical bond formed by sharing a pair of electrons with two atoms.

The explanation of the chemical bond by the appearance of common electron pairs between atoms formed the basis of the spin theory of valence, the tool of which is valence bond method (MVS) , discovered by Lewis in 1916. For the quantum mechanical description of the chemical bond and the structure of molecules, another method is used - molecular orbital method (MMO) .

Valence bond method

The basic principles of the formation of a chemical bond according to MVS:

1. A chemical bond is formed due to valence (unpaired) electrons.

2. Electrons with antiparallel spins belonging to two different atoms become common.

3. A chemical bond is formed only if, when two or more atoms approach each other, the total energy of the system decreases.

4. The main forces acting in the molecule are of electrical, Coulomb origin.

5. The stronger the connection, the more the interacting electron clouds overlap.

There are two mechanisms for the formation of a covalent bond:

exchange mechanism. The bond is formed by sharing the valence electrons of two neutral atoms. Each atom gives one unpaired electron to a common electron pair:

Rice. 7. Exchange mechanism for the formation of a covalent bond: a- non-polar; b- polar

Donor-acceptor mechanism. One atom (donor) provides an electron pair, and another atom (acceptor) provides an empty orbital for this pair.

connections, educated according to the donor-acceptor mechanism, belong to complex compounds

Rice. 8. Donor-acceptor mechanism of covalent bond formation

A covalent bond has certain characteristics.

Saturability - the property of atoms to form a strictly defined number of covalent bonds. Due to the saturation of the bonds, the molecules have a certain composition.

Orientation - t . e. the connection is formed in the direction of maximum overlap of electron clouds . With respect to the line connecting the centers of atoms forming a bond, there are: σ and π (Fig. 9): σ-bond - formed by overlapping AO along the line connecting the centers of interacting atoms; A π-bond is a bond that occurs in the direction of an axis perpendicular to the straight line connecting the nuclei of an atom. The orientation of the bond determines the spatial structure of the molecules, i.e., their geometric shape. hybridization - it is a change in the shape of some orbitals in the formation of a covalent bond in order to achieve a more efficient overlap of orbitals. The chemical bond formed with the participation of electrons of hybrid orbitals is stronger than the bond with the participation of electrons of non-hybrid s- and p-orbitals, since there is more overlap. There are the following types of hybridization (Fig. 10, Table 31): sp hybridization - one s-orbital and one p-orbital turn into two identical "hybrid" orbitals, the angle between the axes of which is 180°. Molecules in which sp hybridization occurs have a linear geometry (BeCl 2).

sp 2 hybridization- one s-orbital and two p-orbitals turn into three identical "hybrid" orbitals, the angle between the axes of which is 120°. Molecules in which sp 2 hybridization is carried out have a flat geometry (BF 3 , AlCl 3).

sp 3-hybridization- one s-orbital and three p-orbitals turn into four identical "hybrid" orbitals, the angle between the axes of which is 109 ° 28 ". Molecules in which sp 3 hybridization occurs have a tetrahedral geometry (CH 4 , NH3).

Rice. 10. Types of hybridizations of valence orbitals: a - sp-hybridization of valence orbitals; b - sp2- hybridization of valence orbitals; in - sp 3 - hybridization of valence orbitals

Algorithm for compiling 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. By 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 has a maximum of 2 electrons. 1s2, 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.
• Filled in first s- sublevel, then p-, d-b f- sublevels.
• Electrons fill orbitals in ascending order of orbital energy (Klechkovsky's rule).
• Within the sublevel, electrons first occupy free orbitals one at a time, and only after that they form pairs (Hund's rule).
• There cannot be more than two electrons in one orbital (Pauli principle).

Examples.

1. Compose the electronic formula of nitrogen. Nitrogen is number 7 on the periodic table.

2. Compose the electronic formula of argon. In the periodic table, argon is at number 18.

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

3. Compose the electronic formula of chromium. In the periodic table, chromium is number 24.

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

Energy diagram of zinc.

4. Compose the electronic formula of zinc. In the periodic table, zinc is number 30.

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

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.