The electronic configuration of chemical elements is the tracking of the location of electrons in its atoms. Electrons can be in shells, subshells, and orbitals. The valency of an element, its chemical activity and the ability to interact with other substances depend on the distribution of electrons.

How electronic configuration is written

The arrangement of atoms is usually written for those particles of chemical elements that are in the ground state. If the atom is excited, the entry will be called the excited configuration. The determination of the electronic configuration applicable in a particular case depends on three rules that are valid for atoms of all chemical elements.

Filling principle

The electronic configuration of an atom must comply with the principle of filling, according to which the electrons of atoms fill the orbits in ascending order - from the lowest energy level to the highest. The lowest orbitals of any atom are always filled first. Then the electrons fill the existing orbitals of the second energy level, then the s orbital, and only at the end - the p-sublevel orbital.

In a letter, the electronic configuration of chemical elements is transmitted by a formula in which, next to the name of the element, a combination of numbers and letters is indicated corresponding to the position of the electrons. The top number indicates the number of electrons in these orbitals.

For example, a hydrogen atom has a single electron. According to the filling principle, this electron is in the s-orbital. Thus, the electronic configuration of hydrogen will be equal to 1s1.

Pauli exclusion principle

The second orbital filling rule is a special case of a more generalized law, which was discovered by the Swiss physicist F. Pauli. According to this rule, in any chemical element there is no pair of electrons that have the same set of quantum numbers. Therefore, no more than two electrons can be in any orbital at the same time, and then only if they have unequal spins.

The Pauli exclusion principle can be seen in specific example. The electronic configuration of a beryllium atom can be written as 1s 2 2s 2 . When an energy quantum hits an atom, the atom goes into an excited state. It can be written like this:

1s 2 2s 2 (normal state) + hν→ 1s 2 2s 1 2p 1 (excited state).

If we compare the electronic configurations of beryllium in the normal and excited states, we can see that the number of unpaired electrons is not the same for them. The electronic configuration of beryllium shows the absence of unpaired electrons in normal state. After a quantum of energy enters the atom, two unpaired electrons appear.

In principle, in any chemical element, electrons can transfer to orbitals with higher energies, but for chemistry, only those transitions that occur between sublevels with close energy values ​​are of interest.

This pattern can be explained as follows. The formation of a chemical bond is always accompanied by the release of energy, because the atoms pass into an energetically favorable state. The depairing of electrons at the same energy level entails such energy costs that are fully compensated after the formation of a chemical bond. The energy costs for the depairing of electrons of different chemical levels turn out to be so great that the chemical bond is not able to compensate for them. If there is no chemical partner, the excited atom releases a quantum of energy and returns to normal condition This process scientists call relaxation.

Gund's rule

The electronic configuration of an atom obeys Hund's law, according to which the filling of the orbitals of one subshell begins with electrons having the same spin. Only after all single electrons occupy the established orbits, charged particles with the opposite spin join them.

Hund's rule clearly confirms the electronic configuration of nitrogen. The nitrogen atom has 7 electrons. The electronic configuration of this chemical element looks like this: ls22s22p3. All three electrons that are located on the 2p subshell must be located one by one, occupying each of the three 2p orbitals, and all of their spins must be parallel.

These rules help not only to understand what determines the electronic configuration of elements periodic system but also to understand the processes occurring inside atoms.

Electronic configuration of an atom is a formula showing the arrangement of electrons in an atom by levels and sublevels. After studying the article, you will find out where and how electrons are located, get acquainted with quantum numbers and be able to build the electronic configuration of an atom by its number, at the end of the article there is a table of elements.

Why study the electronic configuration of elements?

Atoms are like a constructor: there are a certain number of parts, they differ from each other, but two parts of the same type are exactly the same. But this constructor is much more interesting than the plastic one, and here's why. The configuration changes depending on who is nearby. For example, oxygen next to hydrogen Maybe turn into water, next to sodium into gas, and being next to iron completely turns it into rust. To answer the question why this happens and to predict the behavior of an atom next to another, it is necessary to study the electronic configuration, which will be discussed below.

How many electrons are in an atom?

An atom consists of a nucleus and electrons revolving around it, the nucleus consists of protons and neutrons. In the neutral state, each atom has the same number of electrons as the number of protons in its nucleus. The number of protons was indicated by the element's serial number, for example, sulfur has 16 protons - the 16th element of the periodic system. Gold has 79 protons - the 79th element of the periodic table. Accordingly, there are 16 electrons in sulfur in the neutral state, and 79 electrons in gold.

Where to look for an electron?

Observing the behavior of an electron, certain patterns were derived, they are described by quantum numbers, there are four of them in total:

• Principal quantum number
• Orbital quantum number
• Magnetic quantum number
• Spin quantum number

Orbital

Further, instead of the word orbit, we will use the term "orbital", the orbital is the wave function of the electron, roughly - this is the area in which the electron spends 90% of the time.

N - level
L - shell
M l - orbital number
M s - the first or second electron in the orbital

Orbital quantum number l

As a result of the study of the electron cloud, it was found that depending on the level of energy, the cloud takes four main forms: a ball, dumbbells and the other two, more complex. In ascending order of energy, these forms are called s-, p-, d- and f-shells. Each of these shells can have 1 (on s), 3 (on p), 5 (on d) and 7 (on f) orbitals. The orbital quantum number is the shell on which the orbitals are located. The orbital quantum number for s, p, d and f orbitals, respectively, takes the values ​​0,1,2 or 3.

On the s-shell one orbital (L=0) - two electrons
There are three orbitals on the p-shell (L=1) - six electrons
There are five orbitals on the d-shell (L=2) - ten electrons
There are seven orbitals (L=3) on the f-shell - fourteen electrons

Magnetic quantum number m l

There are three orbitals on the p-shell, they are denoted by numbers from -L to +L, that is, for the p-shell (L=1) there are orbitals "-1", "0" and "1". The magnetic quantum number is denoted by the letter m l .

Inside the shell, it is easier for electrons to be located in different orbitals, so the first electrons fill one for each orbital, and then its pair is added to each.

Consider a d-shell:
The d-shell corresponds to the value L=2, that is, five orbitals (-2,-1,0,1 and 2), the first five electrons fill the shell, taking the values ​​M l =-2,M l =-1,M l =0 , M l =1,M l =2.

Spin quantum number m s

Spin is the direction of rotation of an electron around its axis, there are two directions, so the spin quantum number has two values: +1/2 and -1/2. Only two electrons with opposite spins can be on the same energy sublevel. The spin quantum number is denoted m s

Principal quantum number n

The principal quantum number is the level of energy at which this moment seven energy levels are known, each is indicated by an Arabic numeral: 1,2,3, ... 7. The number of shells at each level is equal to the level number: there is one shell on the first level, two on the second, and so on.

Electron number

So, any electron can be described by four quantum numbers, the combination of these numbers is unique for each position of the electron, let's take the first electron, the lowest energy level is N=1, one shell is located on the first level, the first shell at any level has the shape of a ball (s -shell), i.e. L=0, the magnetic quantum number can take only one value, M l =0 and the spin will be equal to +1/2. If we take the fifth electron (in whatever atom it is), then the main quantum numbers for it will be: N=2, L=1, M=-1, spin 1/2.

Task 1. Write the electronic configurations of the following elements: N, Si, F e, Kr , Te, W .

Solution. The energy of atomic orbitals increases in the following order:

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

On each s-shell (one orbital) there can be no more than two electrons, on the p-shell (three orbitals) - no more than six, on the d-shell (five orbitals) - no more than 10 and on the f-shell (seven orbitals) - no more than 14.

In the ground state of an atom, electrons occupy orbitals with the lowest energy. The number of electrons is equal to the charge of the nucleus (the atom as a whole is neutral) and the atomic number of the element. For example, a nitrogen atom has 7 electrons, two of which are in 1s orbitals, two are in 2s orbitals, and the remaining three electrons are in 2p orbitals. The electronic configuration of the nitrogen atom:

7 N : 1s 2 2s 2 2p 3 . Electronic configurations other elements:

14 Si: 1s 2 2s 2 2p 6 3s 2 3p 2 ,

26 F e : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 ,

36 K r: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 3p 6 ,

52 Those : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 3p 6 5s 2 4d 10 5p 4 ,

74 Those : 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 3p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 4 .

Task 2. Which inert gas and ions of which elements have the same electronic configuration as the particle resulting from the removal of all valence electrons from the calcium atom?

Solution. The electron shell of the calcium atom has the structure 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 . When two valence electrons are removed, a Ca 2+ ion is formed with the configuration 1s 2 2s 2 2p 6 3s 2 3p 6 . An atom has the same electronic configuration Ar and ions S 2-, Cl -, K +, Sc 3+, etc.

Task 3. Can the electrons of the Al 3+ ion be in the following orbitals: a) 2p; b) 1r; c) 3d?

Solution. Electronic configuration of aluminum atom: 1s 2 2s 2 2p 6 3s 2 3p 1 . The Al 3+ ion is formed upon the removal of three valence electrons from an aluminum atom and has the electronic configuration 1s 2 2s 2 2p 6 .

a) electrons are already in the 2p orbital;

b) in accordance with the restrictions imposed on the quantum number l (l = 0, 1, ... n -1), at n = 1, only the value l = 0 is possible, therefore, the 1p orbital does not exist;

c) electrons can be in the 3d orbital if the ion is in an excited state.

Task 4. Write the electronic configuration of the neon atom in the first excited state.

Solution. The electronic configuration of the neon atom in the ground state is 1s 2 2s 2 2p 6 . The first excited state is obtained by the transition of one electron from the highest occupied orbital (2p) to the lowest free orbital (3s). The electronic configuration of the neon atom in the first excited state is 1s 2 2s 2 2p 5 3s 1 .

Task 5. What is the composition of nuclei of isotopes 12 C and 13 C , 14 N and 15 N ?

Solution. The number of protons in the nucleus is equal to the atomic number of the element and is the same for all isotopes of this element. The number of neutrons is equal to the mass number (indicated to the upper left of the element number) minus the number of protons. Different isotopes of the same element have different numbers of neutrons.

The composition of these nuclei:

12 C: 6p + 6n; 13 C: 6p + 7n; 14 N : 7p + 7n ; 15N: 7p + 8n.

Electronic configurations of atoms

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

Pauli's rule. Two electrons in one atom 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 the sublevel, electrons have definite 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 the sublevel can contain no more than 2(2l+1) electrons. Hence, the maximum number of electrons at 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.

The rule of least energy. Sequential filling of the levels should occur in such a way as to ensure the minimum energy of the atom. Each electron occupies a free orbital with the lowest energy.

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

Graphic form of the Klechkovsky rule.

According to the Klechkovsky rule, the 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, sublevels are written sequentially by levels: 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 Е 3d

A decrease in the energy of sublevels with smaller n and larger l, if they are completely or half filled, 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 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 and 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 .

Gund's rule. The orbitals of a given sublevel are filled in such a way that the total spin is maximum. The orbitals of a given sublevel are first filled by one electron. For example, for the p 2 configuration, the filling p x 1 p y 1 with a total spin s = 1/2 + 1/2 = 1 is preferable (i.e., it has a lower energy) than the filling p x 2 with a total spin s = 1/2 - 1/2 = 0.

- more profitable, ¯ - less profitable.

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

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

a) by levels)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 sublevels. Titanium is in the 4th period, so we write down the sublevels up to 4p: 1s2s2p3s3p3d4s4p and fill them with electrons up to their total number of 22, while not including the unfilled sublevels in the final formula. We receive.

The electronic configuration of an element is a record of the distribution of electrons in its atoms in shells, subshells and orbitals. The 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 an 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 principle of filling, electrons in the ground state of an atom fill the orbits 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 of all possible orbitals it has the lowest energy (see Fig. 1.21). An electron in this s orbital is called an ls electron. Hydrogen in the ground state has an Is1 electronic configuration.

Rule 2: Pauli exclusion principle. According to this principle, no more than two electrons can be in any orbital, 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 take on two electrons. These electrons must have different 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 the same (parallel) spins cannot be in the same orbital:

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

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

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

In table. 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