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 such 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 the 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 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 the atoms of the elements of the fourth period: 1) we denote the conditionally 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?

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 A.

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.

When writing electronic formulas of atoms of elements, energy levels are indicated (values ​​of the main quantum number n in the form of numbers - 1, 2, 3, etc.), energy sublevels (values ​​of the orbital quantum number l in the form of letters s, p, d, f) and the number at the top indicates the number of electrons in a given sublevel.

The first element in the D.I. Mendeleev is hydrogen, therefore, the charge of the nucleus of an atom H equal to 1, the atom has only one electron per s sublevel of the first level. Therefore, the electronic formula of the hydrogen atom is:

The second element is helium, there are two electrons in its atom, therefore the electronic formula of the helium atom is 2 Not 1s 2. The first period includes only two elements, since the first energy level is filled with electrons, which can only be occupied by 2 electrons.

The third element in order - lithium - is already in the second period, therefore, its second energy level begins to be filled with electrons (we talked about this above). The filling of the second level with electrons begins with s-sublevel, so the electronic formula of the lithium atom is 3 Li 1s 2 2s 1 . In the beryllium atom, filling with electrons is completed s- sublevels: 4 Ve 1s 2 2s 2 .

For subsequent elements of the 2nd period, the second energy level continues to be filled with electrons, only now it is filled with electrons R- sublevel: 5 IN 1s 2 2s 2 2R 1 ; 6 WITH 1s 2 2s 2 2R 2 … 10 Ne 1s 2 2s 2 2R 6 .

Neon atom completes filling with electrons R-sublevel, this element ends the second period, it has eight electrons, since s- And R-sublevels can contain only eight electrons.

The elements of the 3rd period have a similar sequence of filling the energy sublevels of the third level with electrons. The electronic formulas of atoms of some elements of this period are:

11 Na 1s 2 2s 2 2R 6 3s 1 ; 12 mg 1s 2 2s 2 2R 6 3s 2 ; 13 Al 1s 2 2s 2 2R 6 3s 2 3p 1 ;

14 Si 1s 2 2s 2 2R 6 3s 2 3p 2 ;…; 18 Ar 1s 2 2s 2 2R 6 3s 2 3p 6 .

The third period, like the second, ends with an element (argon), which completes its filling with electrons R–sublevel, although the third level includes three sublevels ( s, R, d). According to the above order of filling the energy sublevels in accordance with the rules of Klechkovsky, the energy of sublevel 3 d more sublevel 4 energy s, therefore, the potassium atom following the argon and the calcium atom following it is filled with electrons 3 s- sublevel of the fourth level:

19 TO 1s 2 2s 2 2R 6 3s 2 3p 6 4s 1 ; 20 Sa 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 .

Starting from the 21st element - scandium, in the atoms of the elements, sublevel 3 begins to fill with electrons d. The electronic formulas of the atoms of these elements are:

21 sc 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 3d 1 ; 22 Ti 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 3d 2 .

In the atoms of the 24th element (chromium) and the 29th element (copper), a phenomenon is observed called the “breakthrough” or “failure” of an electron: an electron from an external 4 s-sublevel "fails" by 3 d– sublevel, completing its filling by half (for chromium) or completely (for copper), which contributes to greater stability of the atom:

24 Cr 1s 2 2s 2 2R 6 3s 2 3p 6 4s 1 3d 5 (instead of ...4 s 2 3d 4) and

29 Cu 1s 2 2s 2 2R 6 3s 2 3p 6 4s 1 3d 10 (instead of ...4 s 2 3d 9).

Starting from the 31st element - gallium, the filling of the 4th level with electrons continues, now - R– sublevel:

31 Ga 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 3d 10 4p 1 …; 36 Kr 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 3d 10 4p 6 .

This element ends the fourth period, which already includes 18 elements.

A similar order of filling energy sublevels with electrons takes place in the atoms of elements of the 5th period. The first two (rubidium and strontium) are filled s- sublevel of the 5th level, the next ten elements (from yttrium to cadmium) are filled d– sublevel of the 4th level; six elements complete the period (from indium to xenon), in the atoms of which electrons are filled R-sublevel of the outer, fifth level. There are also 18 elements in a period.

For elements of the sixth period, this filling order is violated. At the beginning of the period, as usual, there are two elements, in the atoms of which is filled with electrons s-sublevel of the outer, sixth, level. At the next element - lanthanum - begins to fill with electrons d–sublevel of the previous level, i.e. 5 d. On this filling with electrons 5 d-sublevel stops and the next 14 elements - from cerium to lutetium - begin to fill f- sublevel of the 4th level. These elements are all included in one cell of the table, and below is an expanded series of these elements, called lanthanides.

Starting from the 72nd element - hafnium - to the 80th element - mercury, filling with electrons continues 5 d- sublevel, and the period ends, as usual, with six elements (from thallium to radon), in the atoms of which it is filled with electrons R-sublevel of the outer, sixth, level. This is the largest period, including 32 elements.

In the atoms of the elements of the seventh, incomplete, period, the same order of filling the sublevels is seen, as described above. We allow students to write electronic formulas of atoms of elements of the 5th - 7th periods, taking into account all that has been said above.

Note:in some textbooks, a different order of writing the electronic formulas of the atoms of the elements is allowed: not in the order in which they are filled, but in accordance with the number of electrons given in the table at each energy level. For example, the electronic formula of an arsenic atom may look like: As 1s 2 2s 2 2R 6 3s 2 3p 6 3d 10 4s 2 4p 3 .

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.

V) 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).

Electronic configuration an atom is a numerical representation of its electron orbitals. Electron orbitals are regions of various shapes located around the atomic nucleus, in which it is mathematically probable that an electron will be found. The electronic configuration helps to quickly and easily tell the reader how many electron orbitals an atom has, as well as determine the number of electrons in each orbital. After reading this article, you will master the method of compiling electronic configurations.

Steps

Distribution of electrons using the periodic system of D. I. Mendeleev

Find the atomic number of your atom. Each atom has a certain number of electrons associated with it. Find the symbol for your atom in the periodic table. The atomic number is a positive integer starting from 1 (for hydrogen) and increasing by one for each subsequent atom. The atomic number is the number of protons in an atom, and therefore it is also the number of electrons in an atom with zero charge.

Determine the charge of an atom. Neutral atoms will have the same number of electrons as shown in the periodic table. However, charged atoms will have more or fewer electrons, depending on the magnitude of their charge. If you are working with a charged atom, add or subtract electrons as follows: add one electron for every negative charge and subtract one for every positive charge.

• For example, a sodium atom with a charge of -1 will have an extra electron in addition to its base atomic number of 11. In other words, an atom will have 12 electrons in total.
• If we are talking about a sodium atom with a charge of +1, one electron must be subtracted from the base atomic number 11. So the atom will have 10 electrons.

1. Memorize the basic list of orbitals. As the number of electrons increases in an atom, they fill the various sublevels of the electron shell of the atom according to a certain sequence. Each sublevel of the electron shell, when filled, contains an even number of electrons. There are the following sublevels:

   Understand the electronic configuration record. Electronic configurations are written down in order to clearly reflect the number of electrons in each orbital. Orbitals are written sequentially, with the number of atoms in each orbital written as a superscript to the right of the orbital name. The completed electronic configuration has the form of a sequence of sublevel designations and superscripts.

   • Here, for example, is the simplest electronic configuration: 1s 2 2s 2 2p 6 . This configuration shows that there are two electrons in the 1s sublevel, two electrons in the 2s sublevel, and six electrons in the 2p sublevel. 2 + 2 + 6 = 10 electrons in total. This is the electronic configuration of the neutral neon atom (neon atomic number is 10).

2. Remember the order of the orbitals. Keep in mind that electron orbitals are numbered in ascending order of electron shell number, but arranged in ascending energy order. For example, a filled 4s 2 orbital has less energy (or less mobility) than a partially filled or filled 3d 10, so the 4s orbital is written first. Once you know the order of the orbitals, you can easily fill them in according to the number of electrons in the atom. The order in which the orbitals are filled is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

   • The electronic configuration of an atom in which all orbitals are filled will have the following form: 10 7p 6
   • Note that the above notation, when all orbits are filled, is the electron configuration of the element Uuo (ununoctium) 118, the highest numbered atom in the Periodic Table. Therefore, this electronic configuration contains all currently known electronic sublevels of a neutrally charged atom.

3. Fill in the orbitals according to the number of electrons in your atom. For example, if we want to write down the electronic configuration of a neutral calcium atom, we must start by looking up its atomic number in the periodic table. Its atomic number is 20, so we will write the configuration of an atom with 20 electrons according to the above order.

   • Fill in the orbitals in the above order until you reach the twentieth electron. The first 1s orbital will have two electrons, the 2s orbital will also have two, the 2p orbital will have six, the 3s orbital will have two, the 3p orbital will have 6, and the 4s orbital will have 2 (2 + 2 + 6 +2 +6 + 2 = 20 .) In other words, the electronic configuration of calcium has the form: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 .
   • Note that the orbitals are in ascending order of energy. For example, when you are ready to move to the 4th energy level, then first write down the 4s orbital, and then 3d. After the fourth energy level, you move on to the fifth, where the same order is repeated. This happens only after the third energy level.

4. Use the periodic table as a visual cue. You have probably already noticed that the shape of the periodic table corresponds to the order of electronic sublevels in electronic configurations. For example, atoms in the second column from the left always end in "s 2 ", while atoms on the right edge of the thin middle section always end in "d 10 ", and so on. Use the periodic table as a visual guide to writing configurations - as the order in which you add to the orbitals corresponds to your position in the table. See below:

   • In particular, the two leftmost columns contain atoms whose electronic configurations end in s-orbitals, the right-hand block of the table contains atoms whose configurations end in p-orbitals, and at the bottom of the atoms end in f-orbitals.
   • For example, when you write down the electronic configuration of chlorine, think like this: "This atom is located in the third row (or "period") of the periodic table. It is also located in the fifth group of the orbital block p of the periodic table. Therefore, its electronic configuration will end with. ..3p 5
   • Note that the elements in the d and f orbital regions of the table have energy levels that do not correspond to the period in which they are located. For example, the first row of a block of elements with d-orbitals corresponds to 3d orbitals, although it is located in the 4th period, and the first row of elements with f-orbitals corresponds to the 4f orbital, despite the fact that it is located in the 6th period.

5. Learn the abbreviations for writing long electronic configurations. The atoms on the right side of the periodic table are called noble gases. These elements are chemically very stable. To shorten the process of writing long electronic configurations, simply write in square brackets the chemical symbol for the nearest noble gas with fewer electrons than your atom, and then continue to write the electronic configuration of subsequent orbital levels. See below:

   • To understand this concept, it will be helpful to write an example configuration. Let's write the configuration of zinc (atomic number 30) using the noble gas abbreviation. The complete zinc configuration looks like this: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 . However, we see that 1s 2 2s 2 2p 6 3s 2 3p 6 is the electronic configuration of argon, a noble gas. Simply replace the electronic configuration part of zinc with the chemical symbol for argon in square brackets (.)
   • So, the electronic configuration of zinc, written in abbreviated form, is: 4s 2 3d 10 .
   • Note that if you are writing the electronic configuration of a noble gas, say argon, you cannot write! One must use the abbreviation of the noble gas in front of this element; for argon it will be neon ().

   Using ADOMAH Periodic Table

   1. Master the ADOMAH periodic table. This method of recording the electronic configuration does not require memorization, however, it requires a modified periodic table, since in the traditional periodic table, starting from the fourth period, the period number does not correspond to the electron shell. Find the ADOMAH periodic table, a special type of periodic table designed by scientist Valery Zimmerman. It is easy to find with a short internet search.

      • In the ADOMAH periodic table, the horizontal rows represent groups of elements such as halogens, noble gases, alkali metals, alkaline earth metals, etc. Vertical columns correspond to electronic levels, and so-called "cascades" (diagonal lines connecting blocks s, p, d and f) correspond to periods.
      • Helium is moved to hydrogen, since both of these elements are characterized by a 1s orbital. The period blocks (s,p,d and f) are shown on the right side and the level numbers are given at the bottom. Elements are represented in boxes numbered from 1 to 120. These numbers are the usual atomic numbers, which represent the total number of electrons in a neutral atom.

   2. Find your atom in the ADOMAH table. To write down the electronic configuration of an element, find its symbol in the ADOMAH periodic table and cross out all elements with a higher atomic number. For example, if you need to write down the electronic configuration of erbium (68), cross out all the elements from 69 to 120.

      • Pay attention to the numbers from 1 to 8 at the base of the table. These are the electronic level numbers, or column numbers. Ignore columns that contain only crossed out items. For erbium, columns with numbers 1,2,3,4,5 and 6 remain.

   3. Count the orbital sublevels up to your element. Looking at the block symbols shown to the right of the table (s, p, d, and f) and the column numbers shown at the bottom, ignore the diagonal lines between the blocks and break the columns into block-columns, listing them in order from bottom to top. And again, ignore the blocks in which all the elements are crossed out. Write the column blocks starting from the column number followed by the block symbol, thus: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 6s (for erbium).

      • Please note: The above electronic configuration Er is written in ascending order of the electronic sublevel number. It can also be written in the order in which the orbitals are filled. To do this, follow the cascades from bottom to top, not columns, when you write column blocks: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 12 .

   4. Count the electrons for each electronic sublevel. Count the elements in each column block that have not been crossed out by attaching one electron from each element, and write their number next to the block symbol for each column block as follows: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 12 5s 2 5p 6 6s 2 . In our example, this is the electronic configuration of erbium.

   5. Be aware of incorrect electronic configurations. There are eighteen typical exceptions related to the electronic configurations of atoms in the lowest energy state, also called the ground energy state. They do not obey the general rule only in the last two or three positions occupied by electrons. In this case, the actual electronic configuration assumes that the electrons are in a state of lower energy compared to the standard configuration of the atom. Exception atoms include:

      • Cr(..., 3d5, 4s1); Cu(..., 3d10, 4s1); Nb(..., 4d4, 5s1); Mo(..., 4d5, 5s1); Ru(..., 4d7, 5s1); Rh(..., 4d8, 5s1); Pd(..., 4d10, 5s0); Ag(..., 4d10, 5s1); La(..., 5d1, 6s2); Ce(..., 4f1, 5d1, 6s2); Gd(..., 4f7, 5d1, 6s2); Au(..., 5d10, 6s1); AC(..., 6d1, 7s2); Th(..., 6d2, 7s2); Pa(..., 5f2, 6d1, 7s2); U(..., 5f3, 6d1, 7s2); Np(..., 5f4, 6d1, 7s2) and cm(..., 5f7, 6d1, 7s2).
   • To find the atomic number of an atom when it is written in electronic form, simply add up all the numbers that follow the letters (s, p, d, and f). This only works for neutral atoms, if you're dealing with an ion it won't work - you'll have to add or subtract the number of extra or lost electrons.
   • The number following the letter is a superscript, do not make a mistake in the control.
   • The "stability of a half-filled" sublevel does not exist. This is a simplification. Any stability that pertains to "half-full" sublevels is due to the fact that each orbital is occupied by one electron, so repulsion between electrons is minimized.
   • Each atom tends to a stable state, and the most stable configurations have filled sublevels s and p (s2 and p6). Noble gases have this configuration, so they rarely react and are located on the right in the periodic table. Therefore, if a configuration ends in 3p 4 , then it needs two electrons to reach a stable state (it takes more energy to lose six, including s-level electrons, so four is easier to lose). And if the configuration ends in 4d 3 , then it needs to lose three electrons to reach a stable state. In addition, half-filled sublevels (s1, p3, d5..) are more stable than, for example, p4 or p2; however, s2 and p6 will be even more stable.
   • When you're dealing with an ion, that means the number of protons is not the same as the number of electrons. The charge of the atom in this case will be shown at the top right (usually) of the chemical symbol. Therefore, an antimony atom with a charge of +2 has the electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1 . Note that 5p 3 has changed to 5p 1 . Be careful when the configuration of a neutral atom ends at sublevels other than s and p. When you take electrons, you can only take them from valence orbitals (s and p orbitals). Therefore, if the configuration ends with 4s 2 3d 7 and the atom gets +2 charge, then the configuration will end with 4s 0 3d 7 . Please note that 3d 7 Not changes, instead electrons of the s-orbital are lost.
   • There are conditions when an electron is forced to "move to a higher energy level." When a sublevel lacks one electron to be half or full, take one electron from the nearest s or p sublevel and move it to the sublevel that needs an electron.
   • There are two options for writing an electronic configuration. They can be written in ascending order of the numbers of energy levels or in the order in which the electron orbitals are filled, as was shown above for erbium.
   • You can also write the electronic configuration of an element by writing only the valence configuration, which is the last s and p sublevel. Thus, the valence configuration of antimony will be 5s 2 5p 3 .
   • Ions are not the same. It's much more difficult with them. Skip two levels and follow the same pattern depending on where you started and how high the number of electrons is.