Let's find out how to write the electronic formula of a chemical element. This question is important and relevant, since it gives an idea not only about the structure, but also about the alleged physical and chemical properties of the atom in question.
Compilation rules
In order to compose a graphical and electronic formula of a chemical element, it is necessary to have an idea of the theory of the structure of the atom. To begin with, there are two main components of an atom: the nucleus and the negative electrons. The nucleus includes neutrons, which have no charge, as well as protons, which have a positive charge.
Arguing how to compose and determine the electronic formula of a chemical element, we note that in order to find the number of protons in the nucleus, the periodic system of Mendeleev is required.
The number of an element in order corresponds to the number of protons in its nucleus. The number of the period in which the atom is located characterizes the number of energy layers on which the electrons are located.
To determine the number of neutrons devoid of an electric charge, it is necessary to subtract its serial number (the number of protons) from the relative mass of an atom of an element.
Instruction
In order to understand how to compose the electronic formula of a chemical element, consider the rule for filling sublevels with negative particles, formulated by Klechkovsky.
Depending on the amount of free energy the free orbitals have, a series is drawn up that characterizes the sequence of filling the levels with electrons.
Each orbital contains only two electrons, which are arranged in antiparallel spins.
In order to express the structure of electron shells, graphic formulas are used. What do the electronic formulas of atoms of chemical elements look like? How to make graphic options? These questions are included in the school chemistry course, so we will dwell on them in more detail.
There is a certain matrix (basis) that is used when compiling graphic formulas. The s-orbital is characterized by only one quantum cell, in which two electrons are located opposite to each other. They are indicated graphically by arrows. For the p orbital, three cells are depicted, each also contains two electrons, ten electrons are located on the d orbital, and f is filled with fourteen electrons.
Examples of compiling electronic formulas
Let's continue the conversation about how to compose the electronic formula of a chemical element. For example, you need to make a graphical and electronic formula for the element manganese. First, we determine the position of this element in the periodic system. It has atomic number 25, so there are 25 electrons in an atom. Manganese is an element of the fourth period, therefore, it has four energy levels.
How to write the electronic formula of a chemical element? We write down the sign of the element, as well as its ordinal number. Using the Klechkovsky rule, we distribute electrons over energy levels and sublevels. We sequentially arrange them on the first, second, and third level, inscribing two electrons in each cell.
Then we sum them up, getting 20 pieces. Three levels are completely filled with electrons, and only five electrons remain on the fourth. Considering that each type of orbital has its own energy reserve, we distribute the remaining electrons to the 4s and 3d sublevels. As a result, the finished electron-graphic formula for the manganese atom has the following form:
1s2/2s2, 2p6/3s2, 3p6/4s2, 3d3
Practical value
With the help of electron-graphic formulas, you can clearly see the number of free (unpaired) electrons that determine the valency of a given chemical element.
We offer a generalized algorithm of actions, with the help of which you can compose electronic graphic formulas of any atoms located in the periodic table.
The first step is to determine the number of electrons using the periodic table. The period number indicates the number of energy levels.
Belonging to a certain group is associated with the number of electrons that are in the outer energy level. The levels are subdivided into sublevels, filled in according to the Klechkovsky rule.
Conclusion
In order to determine the valence capabilities of any chemical element located in the periodic table, it is necessary to draw up an electron-graphic formula of its atom. The algorithm given above will allow to cope with the task, to determine the possible chemical and physical properties of the atom.
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 valency 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 valency 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 electronic structures in a number of analogous elements are only similar, but not identical, then 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 enters, 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
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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.
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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
- First, we write down the sign of the chem. element, where below to the left of the sign we indicate its serial number.
- Further, by the number of the period (from which the element) we determine the number of energy levels and draw next to the sign of the chemical element such a number of arcs.
- Then, according to the group number, the number of electrons in the outer level is written under the arc.
- At the 1st level, the maximum possible is 2e, at the second it is already 8, at the third - as many as 18. We begin to put numbers under the corresponding arcs.
- The number of electrons at the penultimate level must be calculated as follows: the number of already affixed electrons is subtracted from the serial number of the element.
- It remains to turn our circuit into an electronic formula:
- We write the chemical element and its serial number. The number shows the number of electrons in the atom.
- We make a formula. To do this, you need to find out the number of energy levels, the basis for determining the number of the period of the element is taken.
- We break the levels into sub-levels.
Below you can see an example of how to correctly compose electronic formulas of chemical elements.
- first we fill in the s-sublevel, and then the p-, d-b f-sublevels;
- according to the Klechkovsky rule, electrons fill orbitals in order of increasing energy of these orbitals;
- according to Hund's rule, electrons within one sublevel occupy free orbitals one at a time, and then form pairs;
- According to the Pauli principle, there are no more than 2 electrons in one orbital.
The electronic formula of a chemical element shows how many electron layers and how many electrons are contained in an atom and how they are distributed over the layers.
To compile the electronic formula of a chemical element, you need to look at the periodic table and use the information obtained for this element. The serial number of the element in the periodic table corresponds to the number of electrons in the atom. The number of electron layers corresponds to the period number, the number of electrons in the last electron layer corresponds to the group number.
It must be remembered that the first layer has a maximum of 2 1s2 electrons, the second - a maximum of 8 (two s and six p: 2s2 2p6), the third - a maximum of 18 (two s, six p, and ten d: 3s2 3p6 3d10).
For example, the electronic formula of carbon: C 1s2 2s2 2p2 (serial number 6, period number 2, group number 4).
Electronic formula of sodium: Na 1s2 2s2 2p6 3s1 (serial number 11, period number 3, group number 1).
To check the correctness of writing an electronic formula, you can look at the site www.alhimikov.net.
Drawing up an electronic formula of chemical elements at first glance may seem like a rather complicated task, but everything will become clear if you adhere to the following scheme:
- write the orbitals first
- we insert numbers in front of the orbitals that indicate the number of the energy level. Do not forget the formula for determining the maximum number of electrons at the energy level: N=2n2
And how to find out the number of energy levels? Just look at the periodic table: this number is equal to the number of the period in which this element is located.
- above the orbital icon we write a number that indicates the number of electrons that are in this orbital.
For example, the electronic formula for scandium would look like this.
The task of compiling the electronic formula of a chemical element is not the easiest.
So, the algorithm for compiling electronic formulas of elements is as follows:
Here are the electronic formulas of some chemical elements:
You need to compose the electronic formulas of chemical elements in this way: you need to look at the number of the element in the periodic table, thus finding out how many electrons it has. Then you need to find out the number of levels, which is equal to the period. Then the sublevels are written and filled in:
First of all, you need to determine the number of atoms according to the periodic table.
To compile an electronic formula, you will need the periodic system of Mendeleev. Find your chemical element there and look at the period - it will be equal to the number of energy levels. The group number will correspond numerically to the number of electrons in the last level. The element number will be quantitatively equal to the number of its electrons. You also clearly need to know that there are a maximum of 2 electrons on the first level, 8 on the second, and 18 on the third.
These are the highlights. In addition, on the Internet (including our website) you can find information with a ready-made electronic formula for each element, so you can check yourself.
Compiling electronic formulas of chemical elements is a very complex process, you can’t do without special tables, and you need to use a whole bunch of formulas. To summarize, you need to go through these steps:
It is necessary to draw up an orbital diagram in which there will be a concept of the difference between electrons from each other. Orbitals and electrons are highlighted in the diagram.
Electrons are filled in levels, from bottom to top and have several sublevels.
So first we find out the total number of electrons of a given atom.
We fill in the formula according to a certain scheme and write it down - this will be the electronic formula.
For example, for Nitrogen, this formula looks like this, first we deal with electrons:
And write down the formula:
To understand the principle of compiling the electronic formula of a chemical element, first you need to determine the total number of electrons in the atom by the number in the periodic table. After that, you need to determine the number of energy levels, taking as a basis the number of the period in which the element is located.
After that, the levels are broken down into sublevels, which are filled with electrons, based on the Principle of Least Energy.
You can check the correctness of your reasoning by looking, for example, here.
By compiling the electronic formula of a chemical element, you can find out how many electrons and electron layers are in a particular atom, as well as the order in which they are distributed among the layers.
To begin with, we determine the serial number of the element according to the periodic table, it corresponds to the number of electrons. The number of electron layers indicates the period number, and the number of electrons in the last layer of the atom corresponds to the group number.
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.
The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons that have opposite (antiparallel) spins (translated from English as “spindle”), that is, they have properties that can be conditionally represented itself as the rotation of an electron around its imaginary axis: clockwise or counterclockwise. This principle is called the Pauli principle.
If there is one electron in the orbital, then it is called unpaired, if there are two, then these are paired electrons, that is, electrons with opposite spins.
Figure 5 shows a diagram of the division of energy levels into sublevels.
The S-orbital, as you already know, is spherical. The electron of the hydrogen atom (s = 1) is located in this orbital and is unpaired. Therefore, its electronic formula or electronic configuration will be written as follows: 1s 1. In electronic formulas, the energy level number is indicated by the number in front of the letter (1 ...), the sublevel (orbital type) is indicated by the Latin letter, and the number that is written to the upper right of the letter (as an exponent) shows the number of electrons in the sublevel.
For a helium atom, He, having two paired electrons in the same s-orbital, this formula is: 1s 2 .
The electron shell of the helium atom is complete and very stable. Helium is a noble gas.
The second energy level (n = 2) has four orbitals: one s and three p. Second-level s-orbital electrons (2s-orbitals) have a higher energy, since they are at a greater distance from the nucleus than 1s-orbital electrons (n = 2).
In general, for every value of n, there is one s-orbital, but with a corresponding amount of electron energy in it and, therefore, with a corresponding diameter, growing as the value of n increases.
The R-orbital is shaped like a dumbbell or a figure eight. All three p-orbitals are located in the atom mutually perpendicularly along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized again that each energy level (electronic layer), starting from n = 2, has three p-orbitals. As the value of n increases, the electrons occupy p-orbitals located at large distances from the nucleus and directed along the x, y, and z axes.
For elements of the second period (n = 2), first one β-orbital is filled, and then three p-orbitals. Electronic formula 1l: 1s 2 2s 1. The electron is weaker bound to the nucleus of the atom, so the lithium atom can easily give it away (as you obviously remember, this process is called oxidation), turning into a Li + ion.
In the beryllium atom Be 0, the fourth electron is also located in the 2s orbital: 1s 2 2s 2 . The two outer electrons of the beryllium atom are easily detached - Be 0 is oxidized to the Be 2+ cation.
At the boron atom, the fifth electron occupies a 2p orbital: 1s 2 2s 2 2p 1. Further, the atoms C, N, O, E are filled with 2p orbitals, which ends with the noble gas neon: 1s 2 2s 2 2p 6.
For the elements of the third period, the Sv- and Sp-orbitals are filled, respectively. Five d-orbitals of the third level remain free:
Sometimes in diagrams depicting the distribution of electrons in atoms, only the number of electrons at each energy level is indicated, that is, they write down the abbreviated electronic formulas of atoms of chemical elements, in contrast to the full electronic formulas given above.
For elements of large periods (fourth and fifth), the first two electrons occupy the 4th and 5th orbitals, respectively: 19 K 2, 8, 8, 1; 38 Sr 2, 8, 18, 8, 2. Starting from the third element of each large period, the next ten electrons will go to the previous 3d and 4d orbitals, respectively (for elements of secondary subgroups): 23 V 2, 8, 11, 2; 26 Tr 2, 8, 14, 2; 40 Zr 2, 8, 18, 10, 2; 43 Tr 2, 8, 18, 13, 2. As a rule, when the previous d-sublevel is filled, the outer (4p- and 5p, respectively) p-sublevel will begin to fill.
For elements of large periods - the sixth and the incomplete seventh - electronic levels and sublevels are filled with electrons, as a rule, as follows: the first two electrons will go to the outer β-sublevel: 56 Ba 2, 8, 18, 18, 8, 2; 87Gr 2, 8, 18, 32, 18, 8, 1; the next one electron (for Na and Ac) to the previous (p-sublevel: 57 La 2, 8, 18, 18, 9, 2 and 89 Ac 2, 8, 18, 32, 18, 9, 2.
Then the next 14 electrons will go to the third energy level from the outside in the 4f and 5f orbitals, respectively, for lanthanides and actinides.
Then the second outside energy level (d-sublevel) will begin to build up again: for elements of secondary subgroups: 73 Ta 2, 8.18, 32.11, 2; 104 Rf 2, 8.18, 32, 32.10, 2 - and, finally, only after the complete filling of the current level with ten electrons will the outer p-sublevel be filled again:
86 Rn 2, 8, 18, 32, 18, 8.
Very often, the structure of the electron shells of atoms is depicted using energy or quantum cells - they write down the so-called graphic electronic formulas. For this record, the following notation is used: each quantum cell is denoted by a cell that corresponds to one orbital; each electron is indicated by an arrow corresponding to the direction of the spin. When writing a graphical electronic formula, two rules should be remembered: the Pauli principle, according to which there can be no more than two electrons in a cell (orbitals, but with antiparallel spins), and F. Hund's rule, according to which electrons occupy free cells (orbitals), are located in they are first one at a time and at the same time have the same spin value, and only then they pair, but the spins in this case, according to the Pauli principle, will already be oppositely directed.
In conclusion, let us once again consider the mapping of the electronic configurations of the atoms of the elements over the periods of the D. I. Mendeleev system. Schemes of the electronic structure of atoms show the distribution of electrons over electronic layers (energy levels). 
In a helium atom, the first electron layer is completed - it has 2 electrons.
Hydrogen and helium are s-elements; these atoms have an s-orbital filled with electrons.
Elements of the second period
For all elements of the second period, the first electron layer is filled and the electrons fill the e- and p-orbitals of the second electron layer in accordance with the principle of least energy (first s-, and then p) and the rules of Pauli and Hund (Table 2).
In the neon atom, the second electron layer is completed - it has 8 electrons.
Table 2 The structure of the electron shells of atoms of elements of the second period
The end of the table. 2 
Li, Be are β-elements.
B, C, N, O, F, Ne are p-elements; these atoms have p-orbitals filled with electrons.
Elements of the third period
For atoms of elements of the third period, the first and second electron layers are completed; therefore, the third electron layer is filled, in which electrons can occupy the 3s, 3p, and 3d sublevels (Table 3).
Table 3 The structure of the electron shells of atoms of elements of the third period
A 3s-electron orbital is completed at the magnesium atom. Na and Mg are s-elements. 
There are 8 electrons in the outer layer (the third electron layer) in the argon atom. As an outer layer, it is complete, but in total, in the third electron layer, as you already know, there can be 18 electrons, which means that the elements of the third period have unfilled 3d orbitals.
All elements from Al to Ar are p-elements. s- and p-elements form the main subgroups in the Periodic system.
A fourth electron layer appears at the potassium and calcium atoms, and the 4s sublevel is filled (Table 4), since it has a lower energy than the 3d sublevel. To simplify the graphical electronic formulas of atoms of the elements of the fourth period: 1) let's conditionally denote the graphical electronic formula of argon as follows:
Ar;
2) we will not depict the sublevels that are not filled for these atoms.
Table 4 The structure of the electron shells of atoms of the elements of the fourth period
K, Ca - s-elements included in the main subgroups. For atoms from Sc to Zn, the 3d sublevel is filled with electrons. These are 3d elements. They are included in the secondary subgroups, they have a pre-external electron layer filled, they are referred to as transition elements.
Pay attention to the structure of the electron shells of chromium and copper atoms. In them, a "failure" of one electron from the 4n- to the 3d sublevel occurs, which is explained by the greater energy stability of the resulting electronic configurations 3d 5 and 3d 10: 
In the zinc atom, the third electron layer is complete - all the 3s, 3p and 3d sublevels are filled in it, in total there are 18 electrons on them.
In the elements following zinc, the fourth electron layer, the 4p sublevel, continues to be filled: Elements from Ga to Kr are p-elements. 
The outer layer (fourth) of the krypton atom is complete and has 8 electrons. But just in the fourth electron layer, as you know, there can be 32 electrons; the 4d and 4f sublevels of the krypton atom still remain unfilled.
The elements of the fifth period are filling the sublevels in the following order: 5s-> 4d -> 5p. And there are also exceptions associated with the "failure" of electrons, in 41 Nb, 42 MO, etc.
In the sixth and seventh periods, elements appear, that is, elements in which the 4f and 5f sublevels of the third outside electronic layer are being filled, respectively.
The 4f elements are called lanthanides.
5f-elements are called actinides.
The order of filling of electronic sublevels in the atoms of elements of the sixth period: 55 Сs and 56 Ва - 6s-elements;
57 La... 6s 2 5d 1 - 5d element; 58 Ce - 71 Lu - 4f elements; 72 Hf - 80 Hg - 5d elements; 81 Tl - 86 Rn - 6p elements. But even here there are elements in which the order of filling of electronic orbitals is “violated”, which, for example, is associated with greater energy stability of half and completely filled f sublevels, that is, nf 7 and nf 14.
Depending on which sublevel of the atom is filled with electrons last, all elements, as you already understood, are divided into four electronic families or blocks (Fig. 7).
1) s-Elements; the β-sublevel of the outer level of the atom is filled with electrons; s-elements include hydrogen, helium and elements of the main subgroups of groups I and II;
2) p-elements; the p-sublevel of the outer level of the atom is filled with electrons; p elements include elements of the main subgroups of III-VIII groups;
3) d-elements; the d-sublevel of the preexternal level of the atom is filled with electrons; d-elements include elements of secondary subgroups of groups I-VIII, that is, elements of intercalated decades of large periods located between s- and p-elements. They are also called transition elements;
4) f-elements, the f-sublevel of the third outside level of the atom is filled with electrons; these include lanthanides and actinides.
1. What would happen if the Pauli principle was not respected?
2. What would happen if Hund's rule was not respected?
3. Make diagrams of the electronic structure, electronic formulas and graphic electronic formulas of atoms of the following chemical elements: Ca, Fe, Zr, Sn, Nb, Hf, Ra.
4. Write the electronic formula for element #110 using the symbol for the corresponding noble gas.
5. What is the “failure” of an electron? Give examples of elements in which this phenomenon is observed, write down their electronic formulas.
6. How is the belonging of a chemical element to one or another electronic family determined?
7. Compare the electronic and graphic electronic formulas of the sulfur atom. What additional information does the last formula contain?
