Electronic configuration of ca. Electronic configurations of atoms of chemical elements

Electronic configurations atoms

Electrons in an atom occupy levels, sublevels and orbitals according to the following rules.

Pauli's rule. In one atom, two electrons cannot have four identical quantum numbers. They must differ by at least one quantum number.

The orbital contains electrons with certain numbers n, l, m l and the electrons on it can differ only in the quantum number m s, which has two values ​​+1/2 and -1/2. Therefore, no more than two electrons can be located in an orbital.

At a sublevel, electrons have certain n and l and differ in the numbers m l and m s. Since m l can take 2l+1 values, and m s - 2 values, then a sublevel can contain no more than 2(2l+1) electrons. Hence, the maximum numbers of electrons in the s-, p-, d-, f-sublevels are 2, 6, 10, 14 electrons, respectively.

Similarly, a level contains no more than 2n 2 electrons and the maximum number of electrons in the first four levels should not exceed 2, 8, 18 and 32 electrons, respectively.

Rule lowest energy. The sequential filling of levels must occur in such a way as to ensure the minimum energy of the atom. Each electron occupies the vacant orbital with the lowest energy.

Klechkovsky's rule. Filling of electronic sublevels is carried out in increasing order of the sum (n+l), and in the case of the same sum (n+l) - in increasing order of the number n.

Graphic form of Klechkovsky's rule.

According to Klechkovsky's rule, sublevels are filled in the following order: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p, 8s, ...

Although the filling of sublevels occurs according to the Klechkovsky rule, in the electronic formula the sublevels are written sequentially by level: 1s, 2s, 2p, 3s, 3p, 3d, 4s, 4p, 4d, 4f, etc. This is due to the fact that the energy of filled levels is determined by the quantum number n: the larger n, the greater the energy and for completely filled levels we have E 3d

A decrease in the energy of sublevels with smaller n and larger l, if they are filled completely or half, leads for a number of atoms to electronic configurations that differ from those predicted by the Klechkovsky rule. So for Cr and Cu we have a distribution at the valence level:

Cr(24e) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1 and Cu(29e) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 , not

Cr(24e) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 4 4s 2 and Cu(29e) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 9 4s 2 .

Hund's rule. The filling of the orbitals of a given sublevel is carried out so that the total spin is maximum. The orbitals of a given sublevel are filled first with one electron at a time. For example, for configuration p 2, the occupation p x 1 p y 1 with total spin s = 1/2 + 1/2 = 1 is preferable (i.e., it corresponds to lower energy) than the occupation p x 2 with total spin s = 1/2 - 1/2 = 0.

- more profitable, ¯ - less profitable.

Electronic configurations of atoms can be written by levels, sublevels, and orbitals. In the latter case, the orbital is usually designated by a quantum cell, and the electrons by arrows, which have one direction or another depending on the value of m s.

For example, the electronic formula P(15e) can be written:

a) by level)2)8)5

b) by sublevels 1s 2 2s 2 2p 6 3s 2 3p 3

c) by orbitals 1s 2 2s 2 2p x 2 2p y 2 2p z 2 3s 2 3p x 1 3p y 1 3p z 1 or

­ ¯ ­ ¯ ­ ¯ ­ ¯ ­ ¯ ­ ¯ ­ ­ ­

Example. Write down the electronic formulas for Ti(22e) and As(33e) by sublevel. Titan is in the 4th period, so we write down the sublevels up to 4p: 1s2s2p3s3p3d4s4p and fill them with electrons until their total number is 22, while we do not include the unfilled sublevels in the final formula. We get it.

Lewis symbol: Electron diagram: A single electron of a hydrogen atom can participate in the formation of only one chemical bond with other atoms: Number of covalent bonds , which an atom forms in a given compound, characterizes it valence . In all compounds, the hydrogen atom is monovalent. Helium Helium, like hydrogen, is an element of the first period. In its only quantum layer it has one s-an orbital containing two electrons with antiparallel spins (lone electron pair). Lewis symbol: Not:. Electronic configuration 1 s 2, its graphical representation: There are no unpaired electrons in the helium atom, there are no free orbitals. His energy level is complete. Atoms with a complete quantum layer cannot form chemical bonds with other atoms. They're called noble or inert gases. Helium is their first representative. SECOND PERIOD Lithium Atoms of all elements second period have two energy levels. The inner quantum layer is the completed energy level of the helium atom. As shown above, its configuration looks like 1 s 2, but the abbreviated notation can also be used to depict it: . In some literary sources it is designated [K] (by the name of the first electron shell). The second lithium quantum layer contains four orbitals (22 = 4): one s and three R. Electronic configuration of lithium atom: 1 s 22s 1 or 2 s 1. Using the last entry, only the electrons of the outer quantum layer (valence electrons) are isolated. The Lewis symbol for lithium is Li. Graphic representation of electronic configuration:
Beryllium Electronic configuration - 2s2. Electronic diagram of the outer quantum layer:
Bor Electronic configuration - 2s22р1. The boron atom can go into an excited state. Electronic diagram of the outer quantum layer:

In an excited state, a boron atom has three unpaired electrons and can form three chemical bonds: ВF3, B2O3. In this case, the boron atom remains with a free orbital, which can participate in the formation of a bond according to the donor-acceptor mechanism. Carbon Electronic configuration - 2s22р2. Electronic diagrams of the outer quantum layer of a carbon atom in the ground and excited states:

An unexcited carbon atom can form two covalent bonds due to electron pairing and one through the donor-acceptor mechanism. An example of such a compound is carbon monoxide (II), which has the formula CO and is called carbon monoxide. Its structure will be discussed in more detail in section 2.1.2. An excited carbon atom is unique: all orbitals of its outer quantum layer are filled with unpaired electrons, i.e. It has the same number of valence orbitals and valence electrons. Its ideal partner is the hydrogen atom, which has one electron in its only orbital. This explains their ability to form hydrocarbons. Having four unpaired electrons, the carbon atom forms four chemical bonds: CH4, CF4, CO2. In molecules of organic compounds, the carbon atom is always in an excited state:
The nitrogen atom cannot be excited because there is no free orbital in its outer quantum layer. It forms three covalent bonds due to electron pairing:
Having two unpaired electrons in the outer layer, the oxygen atom forms two covalent bonds:
Neon Electronic configuration - 2s22р6. Lewis symbol: Electron diagram of the outer quantum layer:

The neon atom has a complete external energy level and does not form chemical bonds with any atoms. This is the second noble gas. THIRD PERIOD Atoms of all elements of the third period have three quantum layers. The electronic configuration of the two internal energy levels can be depicted as . The outer electronic layer contains nine orbitals, which are populated by electrons, obeying general laws. So, for a sodium atom the electronic configuration is: 3s1, for calcium - 3s2 (in an excited state - 3s13р1), for aluminum - 3s23р1 (in an excited state - 3s13р2). Unlike elements of the second period, atoms of elements of groups V – VII of the third period can exist both in the ground and in excited states. Phosphorus Phosphorus is a group 5 element. Its electronic configuration is 3s23р3. Like nitrogen, it has three unpaired electrons in its outermost energy level and forms three covalent bonds. An example is phosphine, which has the formula PH3 (compare with ammonia). But phosphorus, unlike nitrogen, contains free d-orbitals in the outer quantum layer and can go into an excited state - 3s13р3d1:

This gives it the opportunity to form five covalent bonds in compounds such as P2O5 and H3PO4.

Sulfur The ground state electronic configuration is 3s23p4. Electronic diagram:
However, it can be excited by transferring an electron first from R- on d-orbital (first excited state), and then with s- on d-orbital (second excited state):
In the first excited state, the sulfur atom forms four chemical bonds in compounds such as SO2 and H2SO3. The second excited state of the sulfur atom can be depicted using an electron diagram:

This sulfur atom forms six chemical bonds in the compounds SO3 and H2SO4.

1.3.3. Electronic configurations of atoms of large elements periods THE FOURTH PERIOD

The period begins with potassium (19K) electron configuration: 1s22s22p63s23p64s1 or 4s1 and calcium (20Ca): 1s22s22p63s23p64s2 or 4s2. Thus, in accordance with the Klechkovsky rule, after the p-orbitals of Ar, the outer 4s sublevel is filled, which has lower energy, because The 4s orbital penetrates closer to the nucleus; The 3d sublevel remains empty (3d0). Starting from scandium, the orbitals of the 3d sublevel are populated in 10 elements. They're called d-elements.

In accordance with the principle of sequential filling of orbitals, the chromium atom should have an electronic configuration of 4s23d4, but it exhibits an electron “leap”, which consists in the transition of a 4s electron to a 3d orbital that is close in energy (Fig. 11).

It has been experimentally established that atomic states in which the p-, d-, f-orbitals are half filled (p3, d5, f7), completely (p6, d10, f14) or free (p0, d0, f0) have increased stability. Therefore, if an atom lacks one electron before half-completion or completion of a sublevel, its “leap” from a previously filled orbital (in this case, 4s) is observed.

With the exception of Cr and Cu, all elements from Ca to Zn have the same number of electrons in their outer shell - two. This explains the relatively small change in properties in the series of transition metals. However, for the listed elements, both the 4s electrons of the outer and the 3d electrons of the pre-external sublevel are valence electrons (with the exception of the zinc atom, in which the third energy level is completely completed).

	31Ga 4s23d104p1 32Ge 4s23d104p2 33As 4s23d104p3 34Se 4s23d104p4 35Br 4s23d104p5 36Kr 4s23d104p6

The 4d and 4f orbitals remained free, although the fourth period was completed.

FIFTH PERIOD

The sequence of filling the orbitals is the same as in the previous period: first the 5s orbital is filled ( 37Rb 5s1), then 4d and 5p ( 54Xe 5s24d105p6). The 5s and 4d orbitals are even closer in energy, so most 4d elements (Mo, Tc, Ru, Rh, Pd, Ag) experience an electron transition from the 5s to the 4d sublevel.

SIXTH AND SEVENTH PERIODS

Unlike the previous one, the sixth period includes 32 elements. Cesium and barium are 6s elements. The next energetically favorable states are 6p, 4f and 5d. Contrary to Klechkovsky's rule, in lanthanum it is not the 4f but the 5d orbital that is filled ( 57La 6s25d1), however, for the elements following it, the 4f-sublevel is filled ( 58Ce 6s24f2), on which there are fourteen possible electronic states. Atoms from cerium (Ce) to lutetium (Lu) are called lanthanides - these are f-elements. In the series of lanthanides, sometimes an electron “leak” occurs, just as in the series of d-elements. When the 4f-sublevel is completed, the 5d-sublevel (nine elements) continues to be filled and the sixth period, like any other except the first, is completed by six p-elements.

The first two s elements in the seventh period are francium and radium, followed by one 6d element, actinium ( 89Ac 7s26d1). Actinium is followed by fourteen 5f elements - actinides. The actinides should be followed by nine 6d elements and six p elements should complete the period. The seventh period is incomplete.

The considered pattern of the formation of periods of a system by elements and the filling of atomic orbitals with electrons shows the periodic dependence of the electronic structures of atoms on the charge of the nucleus.

Period is a set of elements arranged in order of increasing charges of atomic nuclei and characterized by the same value of the principal quantum number of outer electrons. At the beginning of the period are filled ns -, and at the end - n.p. -orbitals (except for the first period). These elements form eight main (A) subgroups of the periodic system of D.I. Mendeleev.

Main subgroup is a collection chemical elements, located vertically and having the same number of electrons at the outer energy level.

Within the period, with an increase in the charge of the nucleus and an increasing force of attraction of external electrons to it from left to right, the radii of atoms decrease, which in turn causes a weakening of metallic properties and an increase in non-metallic properties. Behind atomic radius take the theoretically calculated distance from the nucleus to the maximum electron density of the outer quantum layer. In groups, from top to bottom, the number of energy levels increases, and, consequently, the atomic radius. At the same time, the metallic properties are enhanced. Important properties of atoms that change periodically depending on the charges of the atomic nuclei also include ionization energy and electron affinity, which will be discussed in section 2.2.

>> 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 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 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.

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 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 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 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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The electronic configuration of an element is a record of the distribution of electrons in its atoms across shells, subshells and orbitals. Electronic configuration is usually written for atoms in their ground state. The electronic configuration of an atom in which one or more electrons are in an excited state is called the excited configuration. To determine the specific electronic configuration of an element in the ground state, the following three rules exist: Rule 1: filling principle. According to the filling principle, electrons in the ground state of an atom fill orbitals in a sequence of increasing orbital energy levels. The lowest energy orbitals are always filled first.

Hydrogen; atomic number = 1; number of electrons = 1

This single electron in the hydrogen atom must occupy the s orbital of the K-shell, since it has the lowest energy of all possible orbitals (see Fig. 1.21). The electron in this s orbital is called an ls electron. Hydrogen in its ground state has an electronic configuration of Is1.

Rule 2: Pauli's exclusion principle. According to this principle, any orbital can contain no more than two electrons, and then only if they have opposite spins (unequal spin numbers).

Lithium; atomic number = 3; number of electrons = 3

The lowest energy orbital is the 1s orbital. It can only accept two electrons. These electrons must have unequal spins. If we denote spin +1/2 with an arrow pointing up, and spin -1/2 with an arrow pointing down, then two electrons with opposite (antiparallel) spins in the same orbital can be schematically represented by the notation (Fig. 1.27)

Two electrons with identical (parallel) spins cannot exist in one orbital:

The third electron in a lithium atom must occupy the orbital next in energy to the lowest orbital, i.e. 2b-orbital. Thus, lithium has an electronic configuration of Is22s1.

Rule 3: Hund's rule. According to this rule, the filling of the orbitals of one subshell begins with single electrons with parallel (equal sign) spins, and only after single electrons occupy all the orbitals can the final filling of the orbitals with pairs of electrons with opposite spins occur.

Nitrogen; atomic number = 7; number of electrons = 7 Nitrogen has an electron configuration of ls22s22p3. The three electrons located on the 2p subshell must be located singly in each of the three 2p orbitals. In this case, all three electrons must have parallel spins (Fig. 1.22).

In table Figure 1.6 shows the electronic configurations of elements with atomic numbers from 1 to 20.

Table 1.6. Ground state electronic configurations for elements with atomic number 1 to 20

The distribution of electrons over various AOs is called electronic configuration of an atom. The lowest energy electronic configuration corresponds to basic state atom, the remaining configurations refer to excited states.

The electronic configuration of an atom is depicted in two ways - in the form of electronic formulas and electron diffraction diagrams. When writing electronic formulas, the principal and orbital quantum numbers are used. The sublevel is designated using the principal quantum number (number) and the orbital quantum number (corresponding letter). The number of electrons in a sublevel is characterized by the superscript. For example, for the ground state of the hydrogen atom the electronic formula is: 1 s 1 .

The structure of electronic levels can be more fully described using electron diffraction diagrams, where the distribution among sublevels is represented in the form of quantum cells. In this case, the orbital is conventionally depicted as a square with a sublevel designation next to it. The sublevels at each level should be slightly offset in height, since their energy is slightly different. Electrons are represented by arrows or ↓ depending on the sign of the spin quantum number. Electron diffraction diagram of a hydrogen atom:

The principle of constructing electronic configurations of multi-electron atoms is to add protons and electrons to the hydrogen atom. The distribution of electrons across energy levels and sublevels is subject to the rules discussed earlier: the principle of least energy, the Pauli principle and Hund's rule.

Taking into account the structure of the electronic configurations of atoms, all known elements, in accordance with the value of the orbital quantum number of the last filled sublevel, can be divided into four groups: s-elements, p-elements, d-elements, f-elements.

In a helium atom He (Z=2) the second electron occupies 1 s-orbital, its electronic formula: 1 s 2. Electron diffraction diagram:

Helium ends the first shortest period Periodic table elements. The electronic configuration of helium is denoted by .

The second period is opened by lithium Li (Z=3), its electronic formula:
Electron diffraction diagram:

The following are simplified electron diffraction diagrams of atoms of elements whose orbitals of the same energy level are located at the same height. Internal, fully filled sublevels are not shown.

After lithium comes beryllium Be (Z=4), in which an additional electron populates 2 s-orbital. Electronic formula of Be: 2 s 2

In the ground state, the next boron electron B (z=5) occupies 2 R-orbital, V:1 s 2 2s 2 2p 1 ; its electron diffraction diagram:

The following five elements have electronic configurations:

C (Z=6): 2 s 2 2p 2 N (Z=7): 2 s 2 2p 3

O (Z=8): 2 s 2 2p 4 F (Z=9): 2 s 2 2p 5

Ne (Z=10): 2 s 2 2p 6

The given electronic configurations are determined by Hund's rule.

The first and second energy levels of neon are completely filled. Let us denote its electronic configuration and will use it in the future for brevity in writing the electronic formulas of atoms of elements.

Sodium Na (Z=11) and Mg (Z=12) open the third period. Outer electrons occupy 3 s-orbital:

Na (Z=11): 3 s 1

Mg (Z=12): 3 s 2

Then, starting with aluminum (Z=13), fill 3 R-sublevel. The third period ends with argon Ar (Z=18):

Al (Z=13): 3 s 2 3p 1

Ar (Z=18): 3 s 2 3p 6

The elements of the third period differ from the elements of the second in that they have free 3 d-orbitals that can participate in the formation of a chemical bond. This explains the valence states exhibited by elements.

In the fourth period, in accordance with the rule ( n+l), potassium K (Z=19) and calcium Ca (Z=20) have 4 electrons s-sublevel, not 3 d. Starting from scandium Sc (Z=21) and ending with zinc Zn (Z=30), filling 3 d-sublevel:

Electronic formulas d-elements can be represented in ionic form: the sublevels are listed in increasing order of the main quantum number, and at a constant n– in order of increasing orbital quantum number. For example, for Zn such an entry would look like this:
Both of these entries are equivalent, but the zinc formula given earlier correctly reflects the order in which the sublevels are filled.

In row 3 d-elements in chromium Cr (Z=24) there is a deviation from the rule ( n+l). In accordance with this rule, the Cr configuration should look like this:
It has been established that its actual configuration is
This effect is sometimes called electron "dip". Such effects are explained by half the increased resistance ( p 3 , d 5 , f 7) and completely ( p 6 , d 10 , f 14) filled sublevels.

Deviations from the rule ( n+l) are also observed in other elements (Table 2). This is due to the fact that as the principal quantum number increases, the differences between the energies of sublevels decrease.

Next comes filling 4 p-sublevel (Ga - Kr). The fourth period contains only 18 elements. Filling 5 occurs in the same way s-, 4d- and 5 p- sublevels of 18 elements of the fifth period. Note that the energy is 5 s- and 4 d-sublevels are very close, and the electron with 5 s-sublevels can easily move to 4 d-sublevel. At 5 s-sublevel Nb, Mo, Tc, Ru, Rh, Ag has only one electron. In ground state 5 s-Pd sublevel is not filled. A “failure” of two electrons is observed.

table 2

Exceptions from ( n+l) – rules for the first 86 elements

	Electronic configuration
	according to the rule ( n+l)	actual
	4s 2 3d 4 4s 2 3d 9 5s 2 4d 3 5s 2 4d 4 5s 2 4d 5 5s 2 4d 6 5s 2 4d 7 5s 2 4d 8 5s 2 4d 9 6s 2 4f 1 5d 0 6s 2 4f 2 5d 0 6s 2 4f 8 5d 0 6s 2 4f 14 5d 7 6s 2 4f 14 5d 8 6s 2 4f 14 5d 9	4s 1 3d 5 4s 1 3d 10 5s 1 4d 4 5s 1 4d 5 5s 1 4d 6 5s 1 4d 7 5s 1 4d 8 5s 0 4d 10 5s 1 4d 10 6s 2 4f 0 5d 1 6s 2 4f 1 5d 1 6s 2 4f 7 5d 1 6s 0 4f 14 5d 9 6s 1 4f 14 5d 9 6s 1 4f 14 5d 10

In the sixth period after filling 6 s-sublevel of cesium Cs (Z=55) and barium Ba (Z=56) the next electron, according to the rule ( n+l), should take 4 f-sublevel. However, in lanthanum La (Z=57) the electron goes to 5 d-sublevel. Half filled (4 f 7) 4f-sublevel has increased stability, so gadolinium has Gd (Z=64), next to europium Eu (Z=63), by 4 f- the sublevel retains the same number of electrons (7), and a new electron arrives at 5 d-sublevel, breaking the rule ( n+l). In terbium Tb (Z=65) the next electron occupies 4 f-sublevel and the electron transitions from 5 d-sublevel (configuration 4 f 9 6s 2). Filling 4 f-sublevel ends at ytterbium Yb (Z=70). The next electron of the lutetium atom Lu occupies 5 d-sublevel. Its electronic configuration differs from that of the lanthanum atom only in that it is completely filled 4 f-sublevel.

Currently, in the Periodic Table of Elements D.I. Mendeleev under scandium Sc and yttrium Y are sometimes located lutetium (and not lanthanum) as the first d-element, and all 14 elements in front of it, including lanthanum, taking into special group lanthanides beyond the Periodic Table of Elements.

The chemical properties of elements are determined mainly by the structure of the outer electronic levels. Change in the number of electrons on the third outside 4 f-sublevel has little effect on the chemical properties of elements. Therefore all 4 f-elements are similar in their properties. Then in the sixth period the filling of 5 occurs d-sublevel (Hf – Hg) and 6 p-sublevel (Tl – Rn).

In the seventh period 7 s-sublevel is filled with francium Fr (Z=87) and radium Ra (Z=88). Sea anemone exhibits a deviation from the rule ( n+l), and the next electron populates 6 d-sublevel, not 5 f. Next comes a group of elements (Th – No) with 5 being filled f-sublevels that form a family actinides. Note that 6 d- and 5 f- sublevels have such close energies that the electronic configuration of actinide atoms often does not obey the rule ( n+l). But in this case the exact configuration value is 5 f T 5d m not so important, since it has a rather weak effect on Chemical properties element.

In lawrencium Lr (Z=103), a new electron arrives at 6 d-sublevel. This element is sometimes placed under lutetium in the Periodic Table. The seventh period is not completed. Elements 104 – 109 are unstable and their properties are little known. Thus, as the charge of the nucleus increases, similar electronic structures of the outer levels are periodically repeated. In this regard, periodic changes in various properties of elements should also be expected.

Periodic change in the properties of atoms of chemical elements

The chemical properties of atoms of elements are manifested by their interaction. The types of configurations of the external energy levels of atoms determine the main features of their chemical behavior.

The characteristics of the atom of each element that determine its behavior in chemical reactions are ionization energy, electron affinity, and electronegativity.

Ionization energy is the energy required to remove and remove an electron from an atom. The lower the ionization energy, the higher the reducing power of the atom. Therefore, ionization energy is a measure of the reducing power of an atom.

The ionization energy required to remove the first electron is called the first ionization energy I 1 . The energy required to remove the second electron is called the second ionization energy I 2, etc. In this case, the following inequality holds

I 1< I 2 < I 3 .

The separation and removal of an electron from a neutral atom occurs more easily than from a charged ion.

The maximum value of ionization energy corresponds to noble gases. Alkali metals have the minimum ionization energy.

Within one period, the ionization energy changes nonmonotonically. Initially, it decreases when moving from s-elements to the first p-elements. Then it increases in subsequent p-elements.

Within one group, as the atomic number of an element increases, the ionization energy decreases, which is due to an increase in the distance between the outer level and the nucleus.

Electron affinity is the energy (denoted by E) that is released when an electron attaches to an atom. By accepting an electron, the atom becomes a negatively charged ion. Electron affinity increases in a period, but, as a rule, decreases in a group.

Halogens have the highest electron affinity. By adding the electron missing to complete the shell, they acquire the complete configuration of a noble gas atom.

Electronegativity is the sum of ionization energy and electron affinity

Electronegativity increases in a period and decreases in a subgroup.

Atoms and ions do not have strictly defined boundaries due to the wave nature of the electron. Therefore, the radii of atoms and ions are determined conventionally.

The greatest increase in the radius of atoms is observed in elements of small periods, in which only the outer energy level is filled, which is typical for s- and p-elements. For d- and f-elements, a smoother increase in radius is observed with increasing nuclear charge.

Within a subgroup, the radius of atoms increases as the number of energy levels increases.