Ba graphical formula for the distribution of elements. Catalog of files on chemistry

The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons having opposite (antiparallel) spins (translated from English as “spindle”), that is, having such properties that can be conventionally imagined 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, has a spherical shape. 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 preceding the letter (1 ...), Latin letter denote a sublevel (type of orbital), 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, which has two paired electrons in one s-orbital, this formula is: 1s 2.

The electron shell of the helium atom is complete and very stable. Helium is a noble gas.

At the second energy level (n = 2) there are four orbitals: one s and three p. The electrons of the s-orbital of the second level (2s-orbitals) have higher energy, since they are at a greater distance from the nucleus than the electrons of the 1s-orbital (n = 2).

In general, for each value of n there is one s orbital, but with a corresponding supply of electron energy on it and, therefore, with a corresponding diameter, growing as the value of n increases.

The R-Orbital has the shape of a dumbbell or a three-dimensional figure eight. All three p-orbitals are located in the atom mutually perpendicular along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized once again that each energy level (electronic layer), starting from n = 2, has three p-orbitals. As the value of n increases, electrons occupy p-orbitals located at large distances from the nucleus and directed along the x, y, z axes.

For elements of the second period (n = 2), first one b-orbital is filled, and then three p-orbitals. Electronic formula 1l: 1s 2 2s 1. The electron is more loosely bound to the nucleus of the atom, so the lithium atom can easily give it up (as you 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 separated - Be 0 is oxidized into the Be 2+ cation.

In the boron atom, the fifth electron occupies the 2p orbital: 1s 2 2s 2 2p 1. Next, the C, N, O, E atoms are filled with 2p orbitals, which ends with the noble gas neon: 1s 2 2s 2 2p 6.

For elements of the third period, the Sv and Sr 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, abbreviated electronic formulas of atoms are written chemical elements, in contrast to the full electronic formulas given above.

The elements long 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 major period, the next ten electrons will enter the previous 3d and 4d orbitals, respectively (for elements of side subgroups): 23 V 2, 8, 11, 2; 26 Tr 2, 8, 14, 2; 40 Zr 2, 8, 18, 10, 2; 43 Tg 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, like this: the first two electrons will go to the outer b-sublevel: 56 Va 2, 8, 18, 18, 8, 2; 87Gg 2, 8, 18, 32, 18, 8, 1; the next one electron (for Na and Ac) to the previous one (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 enter the third outer energy level in the 4f and 5f orbitals of the lanthanides and actinides, respectively.

Then the second external energy level (d-sublevel) will begin to build up again: for elements of side subgroups: 73 Ta 2, 8.18, 32.11, 2; 104 Rf 2, 8.18, 32, 32.10, 2, - and, finally, only after the current level is completely filled 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 electronic shells of atoms is depicted using energy or quantum cells - so-called graphical electronic formulas are written. For this notation, the following notation is used: each quantum cell is designated by a cell that corresponds to one orbital; Each electron is indicated by an arrow corresponding to the spin direction. When writing a graphical electronic formula, you should remember two rules: the Pauli principle, according to which there can be no more than two electrons in a cell (orbital), but with antiparallel spins, and F. Hund’s rule, according to which electrons occupy free cells (orbitals) and are located in At first, they are one at a time and have the same spin value, and only then they pair, but the spins will be oppositely directed according to the Pauli principle.

In conclusion, let us once again consider the display of electronic configurations of atoms of elements according to the periods of the D.I. Mendeleev system. Diagrams of the electronic structure of atoms show the distribution of electrons across electronic layers (energy levels).

In a helium atom, the first electron layer is complete - it has 2 electrons.

Hydrogen and helium are s-elements; the s-orbital of these atoms is filled with electrons.

Elements of the second period

For all elements of the second period, the first electron layer is filled and 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 Pauli and Hund rules (Table 2).

In the neon atom, the second electron layer is complete - it has 8 electrons.

Table 2 Structure of electronic shells of atoms of elements of the second period

End of table. 2

Li, Be are b-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 electronic layers are completed, so the third electronic layer is filled, in which electrons can occupy the 3s, 3p and 3d sublevels (Table 3).

Table 3 Structure of electronic shells of atoms of elements of the third period

The magnesium atom completes its 3s electron orbital. Na and Mg are s-elements.

An argon atom has 8 electrons in its outer layer (third electron layer). 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. The s- and p-elements form the main subgroups in the Periodic Table.

A fourth electron layer appears in the potassium and calcium atoms, and the 4s sublevel is filled (Table 4), since it has lower energy than the 3d sublevel. To simplify graphical electronic formulas of atoms of elements fourth period: 1) let us denote the conventional graphical electronic formula of argon as follows:
Ar;

2) we will not depict sublevels that are not filled in these atoms.

Table 4 Structure of electronic shells of atoms of elements of the fourth period

K, Ca - s-elements included in the main subgroups. In atoms from Sc to Zn, the 3rd sublevel is filled with electrons. These are Zy elements. They are included in secondary subgroups, their outermost electronic layer is filled, and they are classified as transition elements.

Pay attention to the structure of the electronic shells of chromium and copper atoms. In them there is a “failure” of one electron from the 4th to the 3rd sublevel, which is explained by the greater energy stability of the resulting electronic configurations Zd 5 and Zd 10:

In the zinc atom, the third electron layer is complete - all the 3s, 3p and 3d sublevels are filled in it, with a total of 18 electrons.

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 krypton atom has an outer layer (fourth) that is complete and has 8 electrons. But in total in the fourth electron layer, as you know, there can be 32 electrons; the krypton atom still has unfilled 4d and 4f sublevels.

For elements of the fifth period, sublevels are filled in 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.

4f elements are called lanthanides.

5f-elements are called actinides.

The order of filling electronic sublevels in 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 here, too, there are elements in which the order of filling the electron orbitals is “violated,” which, for example, is associated with the 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 b-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 groups III-VIII;

3) d-elements; the d-sublevel of the pre-external level of the atom is filled with electrons; d-elements include elements of secondary subgroups of groups I-VIII, that is, elements of plug-in 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 outer level of the atom is filled with electrons; these include lanthanides and actinides.

1. What would happen if the Pauli principle were not observed?

2. What would happen if Hund's rule were not followed?

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, Pa.

4. Write the electronic formula for element #110 using the appropriate noble gas symbol.

5. What is an electron “dip”? 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 a particular electronic family determined?

7. Compare the electronic and graphical electronic formulas of the sulfur atom. What additional information does the last formula contain?

>> Chemistry: Electronic configurations of atoms of chemical elements

The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons having opposite (antiparallel) spins (translated from English as “spindle”), that is, having such properties that can be conventionally imagined 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, has a spherical shape. 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 number of the energy level is indicated by the number preceding the letter (1 ...), the Latin letter indicates the sublevel (type of orbital), and the number, which 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, which has two paired electrons in one s-orbital, this formula is: 1s 2.

The electron shell of the helium atom is complete and very stable. Helium is a noble gas.

At the second energy level (n = 2) there are four orbitals: one s and three p. The electrons of the s-orbital of the second level (2s-orbitals) have higher energy, since they are at a greater distance from the nucleus than the electrons of the 1s-orbital (n = 2).

In general, for each value of n there is one s orbital, but with a corresponding supply of electron energy on it and, therefore, with a corresponding diameter, growing as the value of n increases.

The p-Orbital has the shape of a dumbbell or a three-dimensional figure eight. All three p-orbitals are located in the atom mutually perpendicular along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized once again that each energy level (electronic layer), starting from n = 2, has three p-orbitals. As the value of n increases, electrons occupy p-orbitals located at large distances from the nucleus and directed along the x, y, z axes.

For elements of the second period (n = 2), first one b-orbital is filled, and then three p-orbitals. Electronic formula 1l: 1s 2 2s 1. The electron is more loosely bound to the nucleus of the atom, so the lithium atom can easily give it up (as you 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 into the Be 2+ cation.

In the boron atom, the fifth electron occupies the 2p orbital: 1s 2 2s 2 2p 1. Next, the C, N, O, E atoms are filled with 2p orbitals, which ends with the noble gas neon: 1s 2 2s 2 2p 6.

For elements of the third period, the Sv and Sr orbitals are filled, respectively. Five d-orbitals of the third level remain free:

11 Na 1s 2 2s 2 Sv1; 17С11в22822р63р5; 18Аг П^Ёр^Зр6.

Sometimes in diagrams depicting the distribution of electrons in atoms, only the number of electrons at each energy level is indicated, that is, abbreviated electronic formulas of atoms of chemical elements are written, 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 major period, the next ten electrons will enter the previous 3d and 4d orbitals, respectively (for elements of side subgroups): 23 V 2, 8, 11, 2; 26 Tr 2, 8, 14, 2; 40 Zr 2, 8, 18, 10, 2; 43 Tg 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, like this: the first two electrons will go to the outer b-sublevel: 56 Va 2, 8, 18, 18, 8, 2; 87Gg 2, 8, 18, 32, 18, 8, 1; the next one electron (for Na and Ac) to the previous one (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 enter the third outer energy level in the 4f and 5f orbitals of the lanthanides and actinides, respectively.

Then the second external energy level (d-sublevel) will begin to build up again: for elements of side subgroups: 73 Ta 2, 8.18, 32.11, 2; 104 Rf 2, 8.18, 32, 32.10, 2, - and, finally, only after the current level is completely filled 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 electronic shells of atoms is depicted using energy or quantum cells - so-called graphical electronic formulas are written. For this notation, the following notation is used: each quantum cell is designated by a cell that corresponds to one orbital; Each electron is indicated by an arrow corresponding to the spin direction. When writing a graphical electronic formula, you should remember two rules: the Pauli principle, according to which there can be no more than two electrons in a cell (orbital), but with antiparallel spins, and F. Hund’s rule, according to which electrons occupy free cells (orbitals) and are located in At first, they are one at a time and have the same spin value, and only then they pair, but the spins will be oppositely directed according to the Pauli principle.

In conclusion, let us once again consider the display of the electronic configurations of atoms of elements according to the periods of the D.I. Mendeleev system. Diagrams of the electronic structure of atoms show the distribution of electrons across electronic layers (energy levels).

In a helium atom, the first electron layer is complete - it has 2 electrons.

Hydrogen and helium are s-elements; the s-orbital of these atoms is filled with electrons.

Elements of the second period

For all elements of the second period, the first electron layer is filled and 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 Pauli and Hund rules (Table 2).

In the neon atom, the second electron layer is complete - it has 8 electrons.

Table 2 Structure of electronic shells of atoms of elements of the second period

End of table. 2

Li, Be - b-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 electronic layers are completed, so the third electronic layer is filled, in which electrons can occupy the 3s, 3p and 3d sublevels (Table 3).

Table 3 Structure of electronic shells of atoms of elements of the third period

The magnesium atom completes its 3s electron orbital. Na and Mg-s-elements.

An argon atom has 8 electrons in its outer layer (third electron layer). 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. The s- and p-elements form the main subgroups in the Periodic Table.

A fourth electron layer appears in the potassium and calcium atoms, and the 4s sublevel is filled (Table 4), since it has lower energy than the 3d sublevel. To simplify the graphical electronic formulas of atoms of elements of the fourth period: 1) let us denote the conventional graphical electronic formula of argon as follows:
Ar;

2) we will not depict sublevels that are not filled in these atoms.

Table 4 Structure of electronic shells of atoms of elements of the fourth period

K, Ca - s-elements included in the main subgroups. In atoms from Sc to Zn, the 3rd sublevel is filled with electrons. These are Zy elements. They are included in secondary subgroups, their outermost electronic layer is filled, and they are classified as transition elements.

Pay attention to the structure of the electronic shells of chromium and copper atoms. In them there is a “failure” of one electron from the 4th to the 3rd sublevel, which is explained by the greater energy stability of the resulting electronic configurations Zd 5 and Zd 10:

In the zinc atom, the third electron layer is complete - all sublevels 3s, 3p and 3d are filled in it, with a total of 18 electrons.

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 krypton atom has an outer layer (fourth) that is complete and has 8 electrons. But in total in the fourth electron layer, as you know, there can be 32 electrons; the krypton atom still has unfilled 4d and 4f sublevels.

For elements of the fifth period, sublevels are filled in 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.

4f elements are called lanthanides.

5f-elements are called actinides.

The order of filling the 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 here, too, there are elements in which the order of filling the electron orbitals is “violated,” which, for example, is associated with the 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 b-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 II Groups I-VIII;

3) d-elements; the d-sublevel of the pre-external level of the atom is filled with electrons; d-elements include elements of secondary subgroups of groups I-VIII, that is, elements of plug-in 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 outer level of the atom is filled with electrons; these include lanthanides and actinides.

1. What would happen if the Pauli principle were not observed?

2. What would happen if Hund's rule were not followed?

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, Pa.

4. Write the electronic formula for element #110 using the appropriate noble gas symbol.

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6.6. Features of the electronic structure of atoms of chromium, copper and some other elements

If you carefully looked at Appendix 4, you probably noticed that for atoms of some elements the sequence of filling orbitals with electrons is disrupted. Sometimes these violations are called “exceptions,” but this is not so - there are no exceptions to the laws of Nature!

The first element with this disorder is chromium. Let's take a closer look at its electronic structure (Fig. 6.16 A). The chromium atom has 4 s-there are not two sublevels, as one would expect, but only one electron. But at 3 d-sublevel has five electrons, but this sublevel is filled after 4 s-sublevel (see Fig. 6.4). To understand why this happens, let's look at what electron clouds are 3 d-sublevel of this atom.

Each of five 3 d-clouds in this case are formed by one electron. As you already know from § 4 of this chapter, the total electron cloud of such five electrons has a spherical shape, or, as they say, spherically symmetrical. According to the nature of the distribution of electron density in different directions, it is similar to 1 s-EO. The energy of the sublevel whose electrons form such a cloud turns out to be less than in the case of a less symmetrical cloud. In this case, the orbital energy is 3 d-sublevel is equal to energy 4 s-orbitals. When symmetry is broken, for example, when a sixth electron appears, the energy of the orbitals is 3 d-the sublevel again becomes greater than energy 4 s-orbitals. Therefore, the manganese atom again has a second electron at 4 s-AO.
The general cloud of any sublevel, filled with electrons either half or completely, has spherical symmetry. The decrease in energy in these cases is general character and does not depend on whether any sublevel is half or completely filled with electrons. And if so, then we must look for the next violation in the atom in whose electron shell the ninth one “arrives” last d-electron. Indeed, the copper atom has 3 d-sublevel has 10 electrons, and 4 s- only one sublevel (Fig. 6.16 b).
The decrease in the energy of the orbitals of a fully or half-filled sublevel causes a number of important chemical phenomena, some of which you will become familiar with.

6.7. Outer and valence electrons, orbitals and sublevels

In chemistry, the properties of isolated atoms, as a rule, are not studied, since almost all atoms, when part of various substances, form chemical bonds. Chemical bonds are formed by the interaction of electron shells of atoms. For all atoms (except hydrogen), not all electrons take part in the formation of chemical bonds: boron has three out of five electrons, carbon has four out of six, and, for example, barium has two out of fifty-six. These "active" electrons are called valence electrons.

Valence electrons are sometimes confused with external electrons, but this is not the same thing.

Electronic clouds of outer electrons have a maximum radius (and a maximum value of the principal quantum number).

It is the outer electrons that take part in the formation of bonds in the first place, if only because when atoms approach each other, the electron clouds formed by these electrons come into contact first of all. But along with them, some electrons can also take part in the formation of a bond. pre-external(penultimate) layer, but only if they have an energy not very different from the energy of the outer electrons. Both electrons of an atom are valence electrons. (In lanthanides and actinides, even some “outer” electrons are valence)
The energy of valence electrons is much greater than the energy of other electrons of the atom, and valence electrons differ significantly less in energy from each other.
Outer electrons are always valence electrons only if the atom can form chemical bonds at all. Thus, both electrons of the helium atom are external, but they cannot be called valence, since the helium atom does not form any chemical bonds at all.
Valence electrons occupy valence orbitals, which in turn form valence sublevels.

As an example, consider an iron atom, the electronic configuration of which is shown in Fig. 6.17. Of the electrons of an iron atom, the maximum principal quantum number ( n= 4) have only two 4 s-electron. Consequently, they are the outer electrons of this atom. The outer orbitals of the iron atom are all orbitals with n= 4, and the outer sublevels are all the sublevels formed by these orbitals, that is, 4 s-, 4p-, 4d- and 4 f-EPU.
Outer electrons are always valence electrons, therefore 4 s-electrons of the iron atom are valence electrons. And if so, then 3 d-electrons with slightly higher energy will also be valence electrons. At the external level of the iron atom, in addition to the filled 4 s-AO there are still 4 free p-, 4d- and 4 f-AO. All of them are external, but only 4 of them are valence R-AO, since the energy of the remaining orbitals is much higher, and the appearance of electrons in these orbitals is not beneficial for the iron atom.

So, the iron atom
external electronic level – fourth,
external sublevels – 4 s-, 4p-, 4d- and 4 f-EPU,
outer orbitals – 4 s-, 4p-, 4d- and 4 f-AO,
outer electrons – two 4 s-electron (4 s 2),
outer electronic layer – fourth,
external electron cloud – 4 s-EO
valence sublevels – 4 s-, 4p-, and 3 d-EPU,
valence orbitals – 4 s-, 4p-, and 3 d-AO,
valence electrons – two 4 s-electron (4 s 2) and six 3 d-electrons (3 d 6).

Valence sublevels can be filled partially or completely with electrons, or they can remain completely free. As the nuclear charge increases, the energy values ​​of all sublevels decrease, but due to the interaction of electrons with each other, the energy of different sublevels decreases at different “speeds.” Energy fully filled d- And f-sublevels decreases so much that they cease to be valence.

As an example, consider the atoms of titanium and arsenic (Fig. 6.18).

In the case of titanium atom 3 d-EPU is only partially filled with electrons, and its energy is greater than energy 4 s-EPU, and 3 d-electrons are valence. The arsenic atom has 3 d-EPU is completely filled with electrons, and its energy is significantly less than the energy of 4 s-EPU, and therefore 3 d-electrons are not valence.
In the examples given, we analyzed valence electronic configuration titanium and arsenic atoms.

The valence electronic configuration of an atom is depicted as valence electron formula, or in the form energy diagram of valence sublevels.

VALENCE ELECTRONS, EXTERNAL ELECTRONS, VALENCE EPU, VALENCE AO, VALENCE ELECTRON CONFIGURATION OF AN ATOM, VALENCE ELECTRON FORMULA, VALENCE SUBLEVELS DIAGRAM.

1. On the energy diagrams you have compiled and in the complete electronic formulas of the atoms Na, Mg, Al, Si, P, S, Cl, Ar, indicate the outer and valence electrons. Write the valence electronic formulas of these atoms. On the energy diagrams, highlight the parts corresponding to the energy diagrams of the valence sublevels.
2. What do the electronic configurations of atoms have in common: a) Li and Na, B and Al, O and S, Ne and Ar; b) Zn and Mg, Sc and Al, Cr and S, Ti and Si; c) H and He, Li and O, K and Kr, Sc and Ga. What are their differences
3. How many valence sublevels are in the electron shell of an atom of each element: a) hydrogen, helium and lithium, b) nitrogen, sodium and sulfur, c) potassium, cobalt and germanium
4. How many valence orbitals are completely filled in the a) boron, b) fluorine, c) sodium atom?
5. How many orbitals with an unpaired electron does an atom have: a) boron, b) fluorine, c) iron
6. How many free outer orbitals does the manganese atom have? How many free valences?
7.For the next lesson, prepare a strip of paper 20 mm wide, divide it into cells (20 × 20 mm), and apply a natural series of elements (from hydrogen to meitnerium) to this strip.
8.In each cell, place the symbol of the element, its atomic number and valence electron formula, as shown in Fig. 6.19 (use Appendix 4).

6.8. Systematization of atoms according to the structure of their electron shells

The systematization of chemical elements is based on the natural series of elements And principle of similarity of electron shells their atoms.
You are already familiar with the natural series of chemical elements. Now let's get acquainted with the principle of similarity of electronic shells.
Considering the valence electronic formulas of atoms in the ERE, it is easy to discover that for some atoms they differ only in the values ​​of the principal quantum number. For example, 1 s 1 for hydrogen, 2 s 1 for lithium, 3 s 1 for sodium, etc. Or 2 s 2 2p 5 for fluorine, 3 s 2 3p 5 for chlorine, 4 s 2 4p 5 for bromine, etc. This means that the outer regions of the clouds of valence electrons of such atoms are very similar in shape and differ only in size (and, of course, electron density). And if so, then the electron clouds of such atoms and the corresponding valence configurations can be called similar. For atoms of different elements with similar electronic configurations we can write general valence electronic formulas: ns 1 in the first case and ns 2 n.p. 5 in the second. As you move through the natural series of elements, you can find other groups of atoms with similar valence configurations.
Thus, atoms with similar valence electron configurations are regularly found in the natural series of elements. This is the principle of similarity of electronic shells.
Let's try to identify the type of this regularity. To do this, we will use the natural series of elements you made.

The ERE begins with hydrogen, the valence electronic formula of which is 1 s 1 . In search of similar valence configurations, we cut the natural series of elements in front of elements with a common valence electronic formula ns 1 (i.e. before lithium, before sodium, etc.). We received the so-called "periods" of the elements. Let's add the resulting “periods” so that they become table rows (see Fig. 6.20). As a result, only atoms in the first two columns of the table will have similar electronic configurations.

Let's try to achieve similarity of valence electronic configurations in other columns of the table. To do this, we cut out from the 6th and 7th periods elements with numbers 58 – 71 and 90 –103 (they fill 4 f- and 5 f-sublevels) and place them under the table. We will move the symbols of the remaining elements horizontally as shown in the figure. After this, the atoms of elements located in the same column of the table will have similar valence configurations, which can be expressed by general valence electronic formulas: ns 1 , ns 2 , ns 2 (n–1)d 1 , ns 2 (n–1)d 2 and so on until ns 2 n.p. 6. All deviations from the general valence formulas are explained by the same reasons as in the case of chromium and copper (see paragraph 6.6).

As you can see, by using the ERE and applying the principle of similarity of electron shells, we were able to systematize chemical elements. Such a system of chemical elements is called natural, since it is based exclusively on the laws of Nature. The table we received (Fig. 6.21) is one of the ways to graphically represent natural system elements and is called long-period table of chemical elements.

PRINCIPLE OF SIMILARITY OF ELECTRON SHELLS, NATURAL SYSTEM OF CHEMICAL ELEMENTS ("PERIODIC" SYSTEM), TABLE OF CHEMICAL ELEMENTS.

6.9. Long period table of chemical elements

Let's take a closer look at the structure of the long-period table of chemical elements.
The rows of this table, as you already know, are called "periods" of elements. Periods are numbered Arabic numerals from 1 to 7. There are only two elements in the first period. The second and third periods, containing eight elements each, are called short periods. The fourth and fifth periods, containing 18 elements each, are called long periods. The sixth and seventh periods, containing 32 elements each, are called extra long periods.
The columns of this table are called groups elements. Group numbers are indicated by Roman numerals with Latin letters A or B.
Elements of some groups have their own common (group) names: elements of group IA (Li, Na, K, Rb, Cs, Fr) – alkaline elements(or alkali metal elements); Group IIA elements (Ca, Sr, Ba and Ra) – alkaline earth elements(or alkaline earth metal elements)(Name " alkali metals"and alkaline earth metals" refer to the simple substances formed by the corresponding elements and should not be used as names for groups of elements); elements of VIA group (O, S, Se, Te, Po) – chalcogens, group VIIA elements (F, Cl, Br, I, At) – halogens, group VIII elements (He, Ne, Ar, Kr, Xe, Rn) – noble gas elements.(The traditional name "noble gases" also refers to simple substances)
The elements with serial numbers 58 – 71 (Ce – Lu) usually placed at the bottom of the table are called lanthanides(“following lanthanum”), and elements with serial numbers 90 – 103 (Th – Lr) – actinides("following sea anemone"). There is a version of the long-period table, in which lanthanides and actinides are not cut out from the ERE, but remain in their places in ultra-long periods. This table is sometimes called ultra-long-period.
The long period table is divided into four block(or sections).
s-Block includes elements of IA and IIA groups with common valence electronic formulas ns 1 and ns 2 (s-elements).
r-Block includes elements from Group IIIA to VIIIA with common valence electronic formulas from ns 2 n.p. 1 to ns 2 n.p. 6 (p-elements).
d-Block includes elements from group IIIB to IIB with common valence electronic formulas from ns 2 (n–1)d 1 to ns 2 (n–1)d 10 (d-elements).
f-Block includes lanthanides and actinides ( f-elements).

Elements s- And p-blocks form A-groups, and elements d-block – B-group of the system of chemical elements. All f-elements are formally included in group IIIB.
The elements of the first period - hydrogen and helium - are s-elements and can be placed in groups IA and IIA. But helium is more often placed in group VIIIA as the element with which the period ends, which fully corresponds to its properties (helium, like all other simple substances formed by the elements of this group, is a noble gas). Hydrogen is often placed in group VIIA, since its properties are much closer to halogens than to alkaline elements.
Each of the periods of the system begins with an element having a valence configuration of atoms ns 1, since it is from these atoms that the formation of the next electronic layer begins, and ends with an element with a valence configuration of atoms ns 2 n.p. 6 (except for the first period). This makes it easy to identify on the energy diagram groups of sublevels filled with electrons in atoms of each period (Fig. 6.22). Do this work with all the sublevels shown in the copy you made of Figure 6.4. The sublevels highlighted in Figure 6.22 (except for completely filled d- And f-sublevels) are valence for atoms of all elements of a given period.
Appearance in periods s-, p-, d- or f-elements fully correspond to the filling sequence s-, p-, d- or f-sublevels with electrons. This feature of the system of elements allows, knowing the period and group into which a given element belongs, to immediately write down its valence electronic formula.

LONG-PERIOD TABLE OF CHEMICAL ELEMENTS, BLOCKS, PERIODS, GROUPS, ALKALINE ELEMENTS, ALKALINE EARTH ELEMENTS, CHALCOGENS, HALOGENS, NOBLE GASE ELEMENTS, LANTANOIDES, ACTINOIDS.
Write down the general valence electronic formulas of atoms of elements of a) IVA and IVB groups, b) IIIA and VIIB groups?
2. What do the electronic configurations of atoms of elements of groups A and B have in common? How are they different?
3. How many groups of elements are included in a) s-block, b) R-block, c) d-block?
4.Continue Figure 30 in the direction of increasing the energy of the sublevels and highlight groups of sublevels filled with electrons in the 4th, 5th and 6th periods.
5. List the valence sublevels of a) calcium, b) phosphorus, c) titanium, d) chlorine, e) sodium atoms. 6. State how s-, p- and d-elements differ from each other.
7.Explain why the membership of an atom in any element is determined by the number of protons in the nucleus, and not by the mass of this atom.
8.For atoms of lithium, aluminum, strontium, selenium, iron and lead, compose valence, full and abbreviated electronic formulas and draw energy diagrams of valence sublevels. 9.Which element atoms correspond to the following valence electronic formulas: 3 s 1 , 4s 1 3d 1 , 2s 2 2 p 6 , 5s 2 5p 2 , 5s 2 4d 2 ?

6.10. Types of electronic formulas of the atom. Algorithm for their compilation

For different purposes, we need to know either the total or valence configuration of an atom. Each of these electron configurations can be represented by either a formula or an energy diagram. That is, full electron configuration of an atom is expressed full electronic formula of an atom, or complete energy diagram of an atom. In its turn, valence electron configuration of an atom is expressed valence(or as it is often called, " short") electronic formula of the atom, or diagram of valence sublevels of an atom(Fig. 6.23).

Previously, we made electronic formulas for atoms using the atomic numbers of the elements. At the same time, we determined the sequence of filling sublevels with electrons according to the energy diagram: 1 s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s and so on. And only by writing down the complete electronic formula could we write down the valence formula.
It is more convenient to write the valence electronic formula of an atom, which is most often used, based on the position of the element in the system of chemical elements, using period-group coordinates.
Let's take a closer look at how this is done for elements s-, p- And d-blocks
For elements s-block valence electronic formula of an atom consists of three symbols. In general, it can be written as follows:

In the first place (in place of the large cell) the period number is placed (equal to the main quantum number of these s-electrons), and on the third (in superscript) - the group number (equal to the number of valence electrons). Taking the magnesium atom (3rd period, group IIA) as an example, we get:

For elements p-block valence electronic formula of an atom consists of six symbols:

Here, in place of the large cells, the period number is also placed (equal to the main quantum number of these s- And p-electrons), and the group number ( equal to the number valence electrons) turns out to be equal to the sum of the superscripts. For the oxygen atom (2nd period, VIA group) we get:

2s 2 2p 4 .

Valence electronic formula of most elements d-block can be written like this:

As in previous cases, here instead of the first cell the period number is put (equal to the main quantum number of these s-electrons). The number in the second cell turns out to be one less, since the main quantum number of these d-electrons. The group number is here too equal to the sum indexes. Example – valence electronic formula of titanium (4th period, IVB group): 4 s 2 3d 2 .

The group number is equal to the sum of the indices for the elements of the VIB group, but, as you remember, in their valence s-sublevel has only one electron, and the general valence electronic formula is ns 1 (n–1)d 5 . Therefore, the valence electronic formula, for example, of molybdenum (5th period) is 5 s 1 4d 5 .
It is also easy to compose the valence electronic formula of any element of the IB group, for example, gold (6th period)>–>6 s 1 5d 10, but in this case you need to remember that d- the electrons of the atoms of the elements of this group still remain valence, and some of them can participate in the formation of chemical bonds.
The general valence electronic formula of atoms of group IIB elements is ns 2 (n – 1)d 10 . Therefore, the valence electronic formula, for example, of a zinc atom is 4 s 2 3d 10 .
The valence electronic formulas of the elements of the first triad (Fe, Co and Ni) also obey general rules. Iron, an element of group VIIIB, has a valence electronic formula of 4 s 2 3d 6. The cobalt atom has one d-electron more (4 s 2 3d 7), and for the nickel atom - by two (4 s 2 3d 8).
Using only these rules for writing valence electronic formulas, it is impossible to compose electronic formulas for the atoms of some d-elements (Nb, Ru, Rh, Pd, Ir, Pt), since in them, due to the desire for highly symmetrical electron shells, the filling of valence sublevels with electrons has some additional features.
Knowing the valence electronic formula, you can write down the full electronic formula of the atom (see below).
Often, instead of cumbersome complete electronic formulas, they write abbreviated electronic formulas atoms. To compile them in the electronic formula, all the electrons of the atom except the valence ones are isolated, their symbols are placed in square brackets, and the part of the electronic formula corresponding to the electronic formula of the atom of the last element of the previous period (the element forming a noble gas) is replaced with the symbol of this atom.

Examples of electronic formulas of different types are given in Table 14.

Table 14. Examples of electronic formulas of atoms

	Electronic formulas
		Abbreviated	Valence
	1s 2 2s 2 2p 3	2s 2 2p 3	2s 2 2p 3
	1s 2 2s 2 2p 6 3s 2 3p 5	3s 2 3p 5	3s 2 3p 5
	1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5	4s 2 3d 5	4s 2 3d 5
	1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 3	4s 2 4p 3	4s 2 4p 3
	1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6	4s 2 4p 6	4s 2 4p 6

Algorithm for compiling electronic formulas of atoms (using the example of the iodine atom)

№ operations	Operation	Result
	Determine the coordinates of the atom in the table of elements.	Period 5, group VIIA
	Write the valence electron formula.	5s 2 5p 5
	Complete the symbols for the inner electrons in the order in which they fill the sublevels.	1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5
	Considering the decrease in energy of fully filled d- And f-sublevels, write down the complete electronic formula.
	Label the valence electrons.	1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 5
	Identify the electron configuration of the preceding noble gas atom.
	Write down the abbreviated electronic formula by combining square brackets All nonvalent electrons.	5s 2 5p 5

Notes
1. For elements of the 2nd and 3rd periods, the third operation (without the fourth) immediately leads to the complete electronic formula.
2. (n – 1)d 10 -Electrons remain valence on the atoms of group IB elements.

COMPLETE ELECTRONIC FORMULA, VALENCE ELECTRONIC FORMULA, SHORTENED ELECTRONIC FORMULA, ALGORITHM FOR COMPILING ELECTRONIC FORMULAS OF ATOMS.
1. Make up the valence electronic formula of an atom of the element a) the second period of the third A group, b) the third period of the second A group, c) the fourth period of the fourth A group.
2.Make abbreviated electronic formulas for the atoms of magnesium, phosphorus, potassium, iron, bromine and argon.

6.11. Short period table of chemical elements

Over the 100-plus years that have passed since the discovery of the natural system of elements, several hundred different tables have been proposed that graphically reflect this system. Of these, in addition to the long-period table, the most widespread is the so-called short-period table of elements by D. I. Mendeleev. A short-period table is obtained from a long-period table if the 4th, 5th, 6th and 7th periods are cut in front of the elements of the IB group, moved apart and the resulting rows are folded in the same way as we previously folded the periods. The result is shown in Figure 6.24.

Lanthanides and actinides are also placed below the main table here.

IN groups This table contains elements whose atoms same number of valence electrons regardless of what orbitals these electrons are in. Thus, the elements chlorine (a typical element forming a non-metal; 3 s 2 3p 5) and manganese (a metal-forming element; 4 s 2 3d 5), not having similar electron shells, fall here into the same seventh group. The need to distinguish such elements forces us to distinguish them in groups subgroups: main– analogues of the A-groups of the long-period table and side– analogues of B-groups. In Figure 34, the symbols of the elements of the main subgroups are shifted to the left, and the symbols of the elements of the secondary subgroups are shifted to the right.
True, this arrangement of elements in the table also has its advantages, because it is the number of valence electrons that primarily determines the valence capabilities of an atom.
The long-period table reflects the laws of the electronic structure of atoms, the similarities and patterns of changes in the properties of simple substances and compounds across groups of elements, the regular changes in a number of physical quantities characterizing atoms, simple substances and compounds throughout the entire system of elements, and much more. The short-period table is less convenient in this regard.

SHORT-PERIOD TABLE, MAIN SUBGROUPS, SIDE SUBGROUPS.
1. Convert the long-period table you constructed from a natural series of elements into a short-period table. Do the reverse conversion.
2. Is it possible to compile a general valence electronic formula for atoms of elements of one group of the short-period table? Why?

6.12. Atomic sizes. Orbital radii

.

The atom has no clear boundaries. What is considered the size of an isolated atom? The nucleus of an atom is surrounded by an electron shell, and the shell consists of electron clouds. The size of the EO is characterized by a radius r eo. All clouds in the outer layer have approximately the same radius. Therefore, the size of an atom can be characterized by this radius. It is called orbital radius of the atom(r 0).

The values ​​of the orbital radii of atoms are given in Appendix 5.
The radius of the EO depends on the charge of the nucleus and on the orbital in which the electron forming this cloud is located. Consequently, the orbital radius of an atom depends on these same characteristics.
Let's consider the electronic shells of hydrogen and helium atoms. In both the hydrogen atom and the helium atom, electrons are located at 1 s-AO, and their clouds would have the same size if the charges of the nuclei of these atoms were the same. But the charge on the nucleus of a helium atom is twice as large as the charge on the nucleus of a hydrogen atom. According to Coulomb's law, the force of attraction acting on each electron of a helium atom is twice the force of attraction of an electron to the nucleus of a hydrogen atom. Therefore, the radius of the helium atom must be much smaller than the radius of the hydrogen atom. This is true: r 0 (He) / r 0 (H) = 0.291 E / 0.529 E 0.55.
The lithium atom has an outer electron at 2 s-AO, that is, forms a cloud of the second layer. Naturally, its radius should be larger. Really: r 0 (Li) = 1.586 E.
The atoms of the remaining elements of the second period have outer electrons (and 2 s, and 2 p) are located in the same second electron layer, and the nuclear charge of these atoms increases with increasing atomic number. Electrons are more strongly attracted to the nucleus, and, naturally, the radii of the atoms decrease. We could repeat these arguments for atoms of elements of other periods, but with one clarification: the orbital radius decreases monotonically only when each of the sublevels is filled.
But if we ignore the details, the general nature of the change in the sizes of atoms in a system of elements is as follows: with an increase in the ordinal number in a period, the orbital radii of atoms decrease, and in a group they increase. The largest atom is a cesium atom, and the smallest is a helium atom, but of the atoms of elements that form chemical compounds (helium and neon do not form them), the smallest is a fluorine atom.
Most atoms of elements in the natural series after the lanthanides have orbital radii that are somewhat smaller than would be expected based on general laws. This is due to the fact that between lanthanum and hafnium in the system of elements there are 14 lanthanides, and, therefore, the charge of the nucleus of the hafnium atom is 14 e more than lanthanum. Therefore, the outer electrons of these atoms are attracted to the nucleus more strongly than they would be in the absence of lanthanides (this effect is often called “lanthanide contraction”).
Please note that when moving from atoms of group VIIIA elements to atoms of group IA elements, the orbital radius increases abruptly. Consequently, our choice of the first elements of each period (see § 7) turned out to be correct.

ORBITAL RADIUS OF AN ATOM, ITS CHANGE IN THE SYSTEM OF ELEMENTS.
1.According to the data given in Appendix 5, draw on graph paper a graph of the dependence of the orbital radius of an atom on the atomic number of the element for elements with Z from 1 to 40. The length of the horizontal axis is 200 mm, the length of the vertical axis is 100 mm.
2. How can you characterize the appearance of the resulting broken line?

6.13. Atomic ionization energy

If you give an electron in an atom additional energy (you will learn how this can be done in a physics course), then the electron can move to another AO, that is, the atom will end up in excited state. This state is unstable, and the electron will almost immediately return to its original state, and excess energy will be released. But if the energy imparted to the electron is large enough, the electron can completely break away from the atom, while the atom ionized, that is, turns into a positively charged ion ( cation). The energy required for this is called atomic ionization energy(E And).

It is quite difficult to remove an electron from a single atom and measure the energy required for this, so it is practically determined and used molar ionization energy(E and m).

Molar ionization energy shows what is the minimum energy required to remove 1 mole of electrons from 1 mole of atoms (one electron from each atom). This value is usually measured in kilojoules per mole. The values ​​of the molar ionization energy of the first electron for most elements are given in Appendix 6.
How does the ionization energy of an atom depend on the position of the element in the system of elements, that is, how does it change in the group and period?
In its physical meaning, ionization energy is equal to the work that must be expended to overcome the force of attraction between an electron and an atom when moving an electron from an atom to an infinite distance from it.

Where q– electron charge, Q is the charge of the cation remaining after the removal of an electron, and r o is the orbital radius of the atom.

AND q, And Q– the quantities are constant, and we can conclude that the work of removing an electron A, and with it the ionization energy E and, are inversely proportional to the orbital radius of the atom.
By analyzing the values ​​of the orbital radii of atoms of various elements and the corresponding ionization energy values ​​given in Appendices 5 and 6, you can make sure that the relationship between these quantities is close to proportional, but differs somewhat from it. The reason that our conclusion does not agree very well with the experimental data is that we used a very crude model that did not take into account many important factors. But even this rough model allowed us to draw the correct conclusion that with increasing orbital radius the ionization energy of the atom decreases and, conversely, with decreasing radius it increases.
Since in a period with increasing atomic number the orbital radius of atoms decreases, the ionization energy increases. In a group, as the atomic number increases, the orbital radius of atoms, as a rule, increases, and the ionization energy decreases. The highest molar ionization energy is found in the smallest atoms, helium atoms (2372 kJ/mol), and of the atoms capable of forming chemical bonds, fluorine atoms (1681 kJ/mol). The smallest is for the largest atoms, cesium atoms (376 kJ/mol). In a system of elements, the direction of increasing ionization energy can be shown schematically as follows:

In chemistry, it is important that ionization energy characterizes the tendency of an atom to give up “its” electrons: the higher the ionization energy, the less inclined the atom is to give up electrons, and vice versa.

EXCITED STATE, IONIZATION, CATION, IONIZATION ENERGY, MOLAR IONIZATION ENERGY, CHANGE IN IONIZATION ENERGY IN A SYSTEM OF ELEMENTS.
1. Using the data given in Appendix 6, determine how much energy must be expended to remove one electron from all sodium atoms with a total mass of 1 g.
2. Using the data given in Appendix 6, determine how many times more energy is needed to remove one electron from all sodium atoms weighing 3 g than from all potassium atoms of the same mass. Why does this ratio differ from the ratio of the molar ionization energies of the same atoms?
3.According to the data given in Appendix 6, plot the dependence of the molar ionization energy on the atomic number for elements with Z from 1 to 40. The dimensions of the graph are the same as in the assignment to the previous paragraph. Check whether this graph corresponds to the choice of “periods” of the system of elements.

6.14. Electron affinity energy

.

The second most important energy characteristic of an atom is electron affinity energy(E With).

In practice, as in the case of ionization energy, the corresponding molar quantity is usually used - molar electron affinity energy().

Molar electron affinity energy shows the energy released when one mole of electrons is added to one mole of neutral atoms (one electron for each atom). Like molar ionization energy, this quantity is also measured in kilojoules per mole.
At first glance, it may seem that energy should not be released in this case, because an atom is a neutral particle, and there are no electrostatic forces of attraction between a neutral atom and a negatively charged electron. On the contrary, approaching an atom, an electron, it would seem, should be repelled by the same negatively charged electrons that form the electron shell. Actually this is not true. Remember if you have ever had to deal with atomic chlorine. Of course not. After all, it exists only at very high temperatures. Even the more stable molecular chlorine practically does not occur in nature; if necessary, it must be obtained using chemical reactions. And you have to deal with sodium chloride (table salt) constantly. After all, table salt is consumed every day by humans with food. And in nature it occurs quite often. But table salt contains chloride ions, that is, chlorine atoms that have added one “extra” electron. One of the reasons why chloride ions are so common is that chlorine atoms have a tendency to gain electrons, that is, when chloride ions are formed from chlorine atoms and electrons, energy is released.
One of the reasons for the release of energy is already known to you - it is associated with an increase in the symmetry of the electron shell of the chlorine atom during the transition to singly charged anion. At the same time, as you remember, energy 3 p-sublevel decreases. There are other more complex reasons.
Due to the fact that the value of electron affinity energy is influenced by several factors, the nature of the change in this quantity in a system of elements is much more complex than the nature of the change in ionization energy. You can be convinced of this by analyzing the table given in Appendix 7. But since the value of this quantity is determined, first of all, by the same electrostatic interaction as the values ​​of ionization energy, then its change in the system of elements (at least in A- groups) in general outline similar to a change in ionization energy, that is, the energy of electron affinity in a group decreases, and in a period it increases. It is maximum for fluorine (328 kJ/mol) and chlorine (349 kJ/mol) atoms. The nature of the change in electron affinity energy in a system of elements resembles the nature of the change in ionization energy, that is, the direction of increase in electron affinity energy can be shown schematically as follows:

2.On the same scale along the horizontal axis as in previous tasks, construct a graph of the dependence of the molar energy of electron affinity on the atomic number for atoms of elements with Z from 1 to 40 using app 7.
3.What physical meaning do negative electron affinity energy values ​​have?
4. Why, of all the atoms of elements of the 2nd period, only beryllium, nitrogen and neon have negative values ​​of the molar energy of electron affinity?

6.15. The tendency of atoms to lose and gain electrons

You already know that the tendency of an atom to give up its own electrons and to add others’ electrons depends on its energy characteristics (ionization energy and electron affinity energy). Which atoms are more inclined to give up their electrons, and which ones are more inclined to accept others?
To answer this question, let us summarize in Table 15 everything that we know about the change in these inclinations in the system of elements.

Table 15. Changes in the propensity of atoms to give up their own electrons and gain foreign electrons

Now let's consider how many electrons an atom can give up.
Firstly, in chemical reactions an atom can only give up valence electrons, since giving up the rest is energetically extremely unfavorable. Secondly, an atom “easily” gives up (if inclined) only the first electron, it gives up the second electron much more difficult (2-3 times), and the third even more difficult (4-5 times). Thus, an atom can donate one, two and, much less frequently, three electrons.
How many electrons can an atom accept?
Firstly, in chemical reactions an atom can only accept electrons into valence sublevels. Secondly, the release of energy occurs only when the first electron is added (and not always). The addition of a second electron is always energetically unfavorable, and even more so with a third. Nevertheless, an atom can add one, two and (extremely rarely) three electrons, as a rule, as much as it lacks to fill its valence sublevels.
The energy costs for the ionization of atoms and the addition of a second or third electron to them are compensated by the energy released during the formation of chemical bonds. 4. How does the electron shell of potassium, calcium and scandium atoms change when they give up their electrons? Give equations for the release of electrons by atoms and abbreviated electronic formulas for atoms and ions.
5. How does the electron shell of chlorine, sulfur and phosphorus atoms change when they add foreign electrons? Give equations for electron gain and abbreviated electronic formulas for atoms and ions.
6.Using Appendix 7, determine what energy will be released when electrons are added to all sodium atoms total mass 1 year
7. Using Appendix 7, determine how much energy is needed to remove “extra” electrons from 0.1 mole of Br– ions?

Electronic configuration an atom is a numerical representation of its electron orbitals. Electron orbitals are regions various shapes, located around the atomic nucleus, in which it is mathematically probable that an electron will be found. Electronic configuration helps 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 drawing up 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 your atom's symbol on the periodic table. The atomic number is a positive integer starting at 1 (for hydrogen) and increasing by one for each subsequent atom. Atomic number is the number of protons in an atom, and therefore it is also the number of electrons of an atom with zero charge.

Determine the charge of an atom. Neutral atoms will have the same number of electrons as shown on the periodic table. However, charged atoms will have more or less 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 each negative charge and subtract one for each positive charge.

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

1. Remember the basic list of orbitals. As the number of electrons in an atom increases, they fill the various sublevels of the atom's electron shell according to a specific sequence. Each sublevel of the electron shell, when filled, contains even number electrons. The following sublevels are available:

   Understand electronic configuration notation. Electron configurations are written to clearly show 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 takes 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 a neutral neon atom (neon's atomic number is 10).

2. Remember the order of the orbitals. Keep in mind that electron orbitals are numbered in order of increasing electron shell number, but arranged in increasing order of energy. For example, a filled 4s 2 orbital has lower energy (or less mobility) than a partially filled or filled 3d 10 orbital, so the 4s orbital is written first. Once you know the order of the orbitals, you can easily fill them according to the number of electrons in the atom. The order of filling the orbitals 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 be as follows: 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 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p 6
   • Note that the above entry, when all orbitals are filled, is the electron configuration of element Uuo (ununoctium) 118, the atom periodic table with the highest number. Therefore, this electronic configuration contains all the currently known electronic sublevels of a neutrally charged atom.

3. Fill the orbitals according to the number of electrons in your atom. For example, if we want to write down the electron 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 the orbitals according to the order above until you reach the twentieth electron. The first 1s orbital will have two electrons, the 2s orbital will also have two, the 2p will have six, the 3s will have two, the 3p will have 6, and the 4s 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 arranged in order of increasing energy. For example, when you are ready to move to the 4th energy level, first write down the 4s orbital, and then 3d. After the fourth energy level, you move 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've probably already noticed that the shape of the periodic table corresponds to the order of the electron sublevels in the electron configurations. For example, the atoms in the second column from the left always end in "s 2", and the atoms on the right edge of the thin middle part always end in "d 10", etc. Use the periodic table as a visual guide to writing configurations - how the order in which you add to the orbitals corresponds to your position in the table. See below:

   • Specifically, the leftmost two columns contain atoms whose electronic configurations end in s orbitals, the right block of the table contains atoms whose configurations end in p orbitals, and the bottom half contains atoms that 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 p orbital block of the periodic table. Therefore, its electronic configuration will end with. ..3p 5
   • Note that elements in the d and f orbital region of the table are characterized by 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 a 4f orbital, despite being in the 6th period.

5. Learn abbreviations for writing long electron configurations. The atoms on the right edge of the periodic table are called noble gases. These elements are chemically very stable. To shorten the process of writing long electron configurations, simply write the chemical symbol of the nearest noble gas with fewer electrons than your atom in square brackets, and then continue writing the electron 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 abbreviation that includes the noble gas. The complete configuration of zinc 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 electron configuration of argon, a noble gas. Simply replace part of the electronic configuration for zinc with the chemical symbol for argon in square brackets (.)
   • So, the electronic configuration of zinc, written in abbreviated form, has the form: 4s 2 3d 10 .
   • Please note that if you are writing the electronic configuration of a noble gas, say argon, you cannot write it! One must use the abbreviation for the noble gas preceding this element; for argon it will be neon ().

   Using the periodic table ADOMAH

   1. Master the periodic table ADOMAH. This method of recording the electronic configuration does not require memorization, but 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 periodic table ADOMAH - a special type of periodic table developed 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 the so-called "cascades" (diagonal lines connecting blocks s,p,d and f) correspond to periods.
      • Helium is moved towards hydrogen because 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 1 to 120. These numbers are ordinary atomic numbers that represent total electrons in a neutral atom.

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

      • Note the numbers 1 through 8 at the bottom of the table. These are numbers of electronic levels, or numbers of columns. Ignore columns that contain only crossed out items. For erbium, columns numbered 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 base, ignore the diagonal lines between the blocks and break the columns into column blocks, listing them in order from bottom to top. Again, ignore blocks that have all the elements crossed out. Write 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 electron configuration of Er is written in ascending order of electron sublevel number. It can also be written in order of filling the orbitals. To do this, follow the cascades from bottom to top, rather than 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 electron sublevel. Count the elements in each column block that have not been crossed out, attaching one electron from each element, and write their number next to the block symbol for each column block thus: 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 relating to the electronic configurations of atoms in the c state. lowest energy, also called the ground energy state. They don't obey 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 with a 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 electron configuration 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 test.
   • There is no "half-full" sublevel stability. This is a simplification. Any stability that is attributed to "half-filled" sublevels is due to the fact that each orbital is occupied by one electron, thus minimizing repulsion between electrons.
   • Each atom tends to a stable state, and the most stable configurations have the s and p sublevels filled (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 (to lose six, including the s-sublevel electrons, requires more energy, so losing four is easier). And if the configuration ends in 4d 3, then to achieve a stable state it needs to lose three electrons. 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 are dealing with an ion, this means that the number of protons is not equal to the number of electrons. The charge of the atom in this case will be depicted at the top right (usually) of the chemical symbol. Therefore, an antimony atom with charge +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 neutral atom configuration ends in sublevels other than s and p. When you take away electrons, you can only take them from the valence orbitals (s and p orbitals). Therefore, if the configuration ends with 4s 2 3d 7 and the atom receives a charge of +2, then the configuration will end with 4s 0 3d 7. Please note that 3d 7 Not changes, electrons from the s orbital are lost instead.
   • There are conditions when an electron is forced to "move to a higher energy level." When a sublevel is one electron short of being half or full, take one electron from the nearest s or p sublevel and move it to the sublevel that needs the electron.
   • There are two options for recording the electronic configuration. They can be written in increasing order of energy level numbers or in the order of filling electron orbitals, as was shown above for erbium.
   • You can also write the electronic configuration of an element by writing only the valence configuration, which represents 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 large the number of electrons is.

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

The first element in the table is D.I. Mendeleev is hydrogen, therefore the charge of the nucleus of the atom N equals 1, an atom has only one electron per s-sublevel of the first level. Therefore, the electronic formula of the hydrogen atom has the form:

The second element is helium; its atom has two electrons, so 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, therefore the electronic formula of the lithium atom is 3 Li 1s 2 2s 1 . The beryllium atom is completed filling with electrons s-sublevel: 4 Ve 1s 2 2s 2 .

In 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 .

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

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

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 is completely filled with electrons R-sublevel, although the third level includes three sublevels ( s, R, d). According to the above order of filling energy sublevels in accordance with Klechkovsky's rules, the energy of sublevel 3 d more sublevel 4 energy s, therefore, the potassium atom next to argon and the calcium atom behind it are 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 Ca 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 .

Starting from the 21st element - scandium, sublevel 3 in the atoms of the elements begins to be filled 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 called “leakage” or “failure” of an electron is observed: an electron from the outer 4 s– sublevel “falls” by 3 d– sublevel, completing filling it halfway (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 occurs in the atoms of elements of the 5th period. For the first two (rubidium and strontium) it is filled s– sublevel of the 5th level, for the next ten elements (from yttrium to cadmium) is filled d– sublevel of the 4th level; The period is completed by six elements (from indium to xenon), the atoms of which are filled with electrons R– sublevel of the external, fifth level. There are also 18 elements in a period.

For elements of the sixth period, this order of filling is violated. At the beginning of the period, as usual, there are two elements whose atoms are filled with electrons s– sublevel of the external, sixth, level. The next element behind them, lanthanum, begins to fill with electrons d– sublevel of the previous level, i.e. 5 d. This completes the 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 row 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), the atoms of which are filled with electrons R– sublevel of the external, 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 sublevels is visible as described above. We let students write the electronic formulas of atoms of elements of the 5th – 7th periods themselves, taking into account everything said above.

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