Inorganic chemistry. Electronic configurations of atoms of elements of the Periodic Table

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:

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

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

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).

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 set of chemical elements arranged 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.

Electronic configuration atom is a formula describing the arrangement of electrons in different electron shells of an atom of a chemical element. The number of electrons in a neutral atom is numerically equal to the charge of the nucleus, and, therefore, the serial number in the periodic table.

As the number of electrons in an atom increases, they fill different sublevels of the atom's electron shell. Each sublevel of the electron shell, when filled, contains an even number of electrons:

- s-sublayer contains a single orbital, which, according to Pauli, can contain a maximum of two electrons.

- p-sublevel contains three orbitals, and can therefore hold a maximum of 6 electrons.

- d-sublevel contains 5 orbitals, so it can have up to 10 electrons.

- f-sublevel contains 7 orbitals, so it can have up to 14 electrons.

Electronic orbitals are numbered in increasing order of the principal quantum number (level number), which coincides with the period number. Orbitals are filled in order of increasing energy (principle of minimum energy): 1 s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.If you know the order of filling the orbitals and understand that each subsequent atom of an element in the periodic table has one more electron than the previous one, it is easy to fill them in accordance with the number of electrons in the atom.

Only electrons of the outer level of the atom—valence electrons—participate in chemical transformations. The elements that complete the periods of the periodic table, the noble gases, which have completely filled electron orbitals, are chemically very stable. To write down a brief electronic configuration of atom A, it is enough to write in square brackets the chemical symbol of the nearest inert gas with a smaller number of electrons than atom A, and then add the configuration of subsequent orbital sublevels.

A graphical representation of the electron configuration demonstrates the arrangement of electrons across quantum cells. Quantum cells should be positioned relative to each other, taking into account the energy of the orbitals. Cells of energetically degenerate orbitals are located at the same level, more energetically favorable ones are lower, less favorable ones are higher. The table shows the electronic configuration of the arsenic atom. Filled as well as half filled d- sublevels have lower orbital energies than s- sublevels, so they are drawn below. Table 2 shows the configuration for the arsenic atom.

Table 2. Electronic configuration of the arsenic atom As

There are exceptions to the electronic configurations of atoms in the ground energy state, for example: Cr (3 d 5 4s 1); Cu (3 d 10 4s 1); Mo (4 d 5 5s 1); Ag (4 d 10 5s 1); Au (4 f 14 5d 10 6s 1 .

Chemical bond

The properties of a substance are determined by its chemical composition, the order in which atoms are combined into molecules and crystal lattices, and their mutual influence. The electronic structure of each atom determines the mechanism of formation of chemical bonds, its type and characteristics.

Initially, the elements in the Periodic Table of Chemical Elements by D.I. Mendeleev's atoms were arranged in accordance with their atomic masses and chemical properties, but in fact it turned out that the decisive role is played not by the mass of the atom, but by the charge of the nucleus and, accordingly, the number of electrons in a neutral atom.

The most stable state of an electron in an atom of a chemical element corresponds to the minimum of its energy, and any other state is called excited, in which the electron can spontaneously move to a level with a lower energy.

Let's consider how electrons in an atom are distributed among orbitals, i.e. electronic configuration of a multielectron atom in the ground state. To construct the electronic configuration, the following principles are used for filling orbitals with electrons:

- Pauli principle (prohibition) - in an atom there cannot be two electrons with the same set of all 4 quantum numbers;

- the principle of least energy (Klechkovsky's rules) - the orbitals are filled with electrons in order of increasing energy of the orbitals (Fig. 1).

Rice. 1. Energy distribution of orbitals of a hydrogen-like atom; n is the principal quantum number.

The energy of the orbital depends on the sum (n + l). The orbitals are filled with electrons in order of increasing sum (n + l) for these orbitals. Thus, for the 3d and 4s sublevels, the sums (n + l) will be equal to 5 and 4, respectively, as a result of which the 4s orbital will be filled first. If the sum (n + l) is the same for two orbitals, then the orbital with the smaller n value is filled first. So, for 3d and 4p orbitals, the sum (n + l) will be equal to 5 for each orbital, but the 3d orbital is filled first. According to these rules, the order of filling the orbitals will be as follows:

1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<5d<4f<6p<7s<6d<5f<7p

An element's family is determined by the last orbital to be filled by electrons, according to energy. However, it is impossible to write electronic formulas in accordance with the energy series.

41 Nb 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 3 5s 2 correct notation of electronic configuration

41 Nb 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 3 incorrect electronic configuration entry

For the first five d - elements, the valence (i.e., electrons responsible for the formation of a chemical bond) is the sum of the electrons on d and s, the last ones filled with electrons. For p-elements, the valence is the sum of the electrons located in the s and p sublevels. For s elements, the valence electrons are the electrons located in the s sublevel of the outer energy level.

- Hund's rule - at one value of l, electrons fill the orbitals in such a way that the total spin is maximum (Fig. 2)

Rice. 2. Change in energy in the 1s -, 2s – 2p – orbitals of atoms of the 2nd period of the Periodic Table.

Examples of constructing electronic configurations of atoms

Examples of constructing electronic configurations of atoms are given in Table 1.

Table 1. Examples of constructing electronic configurations of atoms

>> 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 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.

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