The structure of the electronic shells of atoms. Electronic formula of an element Electronic structure of elements table

s-Elements Elements in the atoms of which the last electron enters the s-sublevel are called. Similarly defined p-elements,d-elements andf-elements.

The beginning of each period corresponds to the opening of a new electronic layer. The period number is equal to the number of the electron layer being opened. Each period, except the first, ends with the filling of the p-sublevel of the layer opened at the beginning of this period. The first period contains only s-elements (two). In the fourth and fifth periods, between the s-elements (two) and p-elements (six) there are d-elements (ten). In the sixth and seventh, behind a pair of s-elements there is (in violation of Klechkovsky's rules) one d-element, then fourteen f-elements (they are placed in separate rows at the bottom of the table - lanthanides and actinides), then nine d-elements and, as always , periods end with six p-elements.

The table is divided vertically into 8 groups, each group into a main and secondary subgroup. The main subgroups contain s- and p-elements, and the secondary subgroups contain d-elements. The main subgroup is easy to determine - it contains elements of periods 1-3. Strictly below them are the remaining elements of the main subgroup. Elements of the side subgroup are located to the side (left or right).

Valence of atoms

In the classical concept, valence is determined by the number of unpaired electrons in the ground or excited state of atoms. Ground state- the electronic state of an atom in which its energy is minimal. Excited state- the electronic state of an atom corresponding to the transition of one or more electrons from an orbital with lower energy to a free orbital with higher energy. For s- and p-elements, electron transition is possible only within the outer electron layer. For d-elements, transitions are possible within the d-sublevel of the pre-external layer and the s- and p-sublevels of the outer layer. For f-elements, transitions are possible within the (n-2)f-, (n-1)d-, ns- and np-sublevels, where n is the number of the outer electronic layer. Valence electrons are called electrons that determine the valence of an atom in its ground or excited state. Valence electron layer- layer on which valence electrons are located.

Describe the electrons of the outer layer of the sulfur atom and the valence electrons of iron (ground state) using quantum numbers. Indicate the possible valencies and oxidation states of the atoms of these elements.

1). Sulfur atom.

Sulfur has serial number 16. It is in the third period, sixth group, main subgroup. Therefore, this is a p-element, the outer electron layer is the third, it is the valence one. It has six electrons. The electronic structure of the valence layer has the form

   

For all electrons n=3, since they are located on the third layer. Let's look at them in order:

 n=3, L=0 (the electron is located in the s-orbital), m l =0 (for the s-orbital only this value of the magnetic quantum number is possible), m s =+1/2 (rotation around its own axis occurs clockwise) ;

 n=3, L=0, m l =0 (these three quantum numbers are the same as those of the first electron, since both electrons are in the same orbital), m s = -1/2 (only here the difference appears, required by the Pauli principle);

 n=3, L=1 (this is a p-electron), m l =+1 (out of three possible values ​​m l = 1, 0 for the first p-orbital we choose the maximum, this is a p x orbital), m s = +1/ 2;

 n=3, L=1, m l = +1, m s =-1/2;

 n=3, L=1, m l = 0 (this is a p y orbital), m s = +1/2;

 n=3, L=1, m l = -1 (this is a p z orbital), m s = +1/2.

Let's consider the valence and oxidation states of sulfur. On the valence layer in the ground state of the atom there are two electron pairs, two unpaired electrons and five free orbitals. Therefore, the valency of sulfur in this state is II. Sulfur is a non-metal. It lacks two electrons before completing the layer, so in compounds with atoms of less electronegative elements, such as metals, it can exhibit a minimum oxidation state of -2. Pairing of electron pairs is possible because there are free orbitals on this layer. Therefore, in the first excited state (S *)

In compounds with atoms of more electronegative elements, such as oxygen, all six valence electrons can be displaced from the sulfur atoms, so its maximum oxidation state is +6.

2). Iron.

The serial number of iron is 26. It is located in the fourth period, in the eighth group, a secondary subgroup. This is a d-element, the sixth in a series of d-elements of the fourth period. Iron valence electrons (eight) are located on the 3d sublevel (six, in accordance with their position in the series of d elements) and on the 4s sublevel (two):

    

Let's look at them in order:

 n=3, L=2, m l = +2, m s = +1/2;

 n=3, L=2, m l = +2, m s = -1/2;

 n=3, L=2, m l = +1, m s = +1/2;

 n=3, L=2, m l = 0, m s = +1/2;

 n=3, L=2, m l = -1, m s = +1/2;

 n=3, L=2, m l = -2, m s = +1/2;

 n=4, L=0, m l = 0, m s = +1/2;

 n=4, L=0, m l = 0, m s = -1/2.

Valence

There are no unpaired electrons on the outer layer, so the minimum valency of iron (II) appears in the excited state of the atom:

After the electrons of the outer layer are used, 4 unpaired electrons of the 3d sublevel can be involved in the formation of chemical bonds. Therefore, the maximum valence of iron is VI.

Oxidation state

Iron is a metal, so it is characterized by positive oxidation states from +2 (electrons of the 4s sublevel are involved) to +6 (4s and all unpaired 3d electrons are involved).

Chemicals are what the world around us is made of.

The properties of each chemical substance are divided into two types: chemical, which characterize its ability to form other substances, and physical, which are objectively observed and can be considered in isolation from chemical transformations. For example, the physical properties of a substance are its state of aggregation (solid, liquid or gaseous), thermal conductivity, heat capacity, solubility in various media (water, alcohol, etc.), density, color, taste, etc.

The transformation of some chemical substances into other substances is called chemical phenomena or chemical reactions. It should be noted that there are also physical phenomena that are obviously accompanied by a change in any physical properties of a substance without its transformation into other substances. Physical phenomena, for example, include the melting of ice, freezing or evaporation of water, etc.

The fact that a chemical phenomenon is taking place during a process can be concluded by observing characteristic signs of chemical reactions, such as color changes, the formation of precipitates, the release of gas, the release of heat and (or) light.

For example, a conclusion about the occurrence of chemical reactions can be made by observing:

Formation of sediment when boiling water, called scale in everyday life;

The release of heat and light when a fire burns;

Change in color of a cut of a fresh apple in air;

Formation of gas bubbles during dough fermentation, etc.

The smallest particles of a substance that undergo virtually no changes during chemical reactions, but only connect with each other in a new way, are called atoms.

The very idea of ​​the existence of such units of matter arose in ancient Greece in the minds of ancient philosophers, which actually explains the origin of the term “atom,” since “atomos” literally translated from Greek means “indivisible.”

However, contrary to the idea of ​​ancient Greek philosophers, atoms are not the absolute minimum of matter, i.e. they themselves have a complex structure.

Each atom consists of so-called subatomic particles - protons, neutrons and electrons, designated respectively by the symbols p +, n o and e -. The superscript in the notation used indicates that the proton has a unit positive charge, the electron has a unit negative charge, and the neutron has no charge.

As for the qualitative structure of an atom, in each atom all protons and neutrons are concentrated in the so-called nucleus, around which the electrons form an electron shell.

The proton and neutron have almost the same masses, i.e. m p ≈ m n, and the mass of the electron is almost 2000 times less than the mass of each of them, i.e. m p /m e ≈ m n /m e ≈ 2000.

Since the fundamental property of an atom is its electrical neutrality, and the charge of one electron is equal to the charge of one proton, from this we can conclude that the number of electrons in any atom is equal to the number of protons.

For example, the table below shows the possible composition of atoms:

Type of atoms with the same nuclear charge, i.e. with the same number of protons in their nuclei is called a chemical element. Thus, from the table above we can conclude that atom1 and atom2 belong to one chemical element, and atom3 and atom4 belong to another chemical element.

Each chemical element has its own name and individual symbol, which is read in a certain way. So, for example, the simplest chemical element, the atoms of which contain only one proton in the nucleus, is called “hydrogen” and is denoted by the symbol “H”, which is read as “ash”, and a chemical element with a nuclear charge of +7 (i.e. containing 7 protons) - “nitrogen”, has the symbol “N”, which is read as “en”.

As you can see from the table above, atoms of one chemical element can differ in the number of neutrons in their nuclei.

Atoms that belong to the same chemical element, but have a different number of neutrons and, as a result, mass, are called isotopes.

For example, the chemical element hydrogen has three isotopes - 1 H, 2 H and 3 H. The indices 1, 2 and 3 above the symbol H mean the total number of neutrons and protons. Those. Knowing that hydrogen is a chemical element, which is characterized by the fact that there is one proton in the nuclei of its atoms, we can conclude that in the 1 H isotope there are no neutrons at all (1-1 = 0), in the 2 H isotope - 1 neutron (2-1=1) and in the 3 H isotope – two neutrons (3-1=2). Since, as already mentioned, the neutron and proton have the same masses, and the mass of the electron is negligibly small in comparison with them, this means that the 2 H isotope is almost twice as heavy as the 1 H isotope, and the 3 H isotope is even three times heavier . Due to such a large scatter in the masses of hydrogen isotopes, the isotopes 2 H and 3 H were even assigned separate individual names and symbols, which is not typical for any other chemical element. The 2H isotope was named deuterium and given the symbol D, and the 3H isotope was given the name tritium and given the symbol T.

If we take the mass of the proton and neutron as one, and neglect the mass of the electron, in fact the upper left index, in addition to the total number of protons and neutrons in the atom, can be considered its mass, and therefore this index is called the mass number and is designated by the symbol A. Since the charge of the nucleus of any Protons correspond to the atom, and the charge of each proton is conventionally considered equal to +1, the number of protons in the nucleus is called the charge number (Z). By denoting the number of neutrons in an atom as N, the relationship between mass number, charge number, and number of neutrons can be expressed mathematically as:

According to modern concepts, the electron has a dual (particle-wave) nature. It has the properties of both a particle and a wave. Like a particle, an electron has mass and charge, but at the same time, the flow of electrons, like a wave, is characterized by the ability to diffraction.

To describe the state of an electron in an atom, the concepts of quantum mechanics are used, according to which the electron does not have a specific trajectory of motion and can be located at any point in space, but with different probabilities.

The region of space around the nucleus where an electron is most likely to be found is called an atomic orbital.

An atomic orbital can have different shapes, sizes, and orientations. An atomic orbital is also called an electron cloud.

Graphically, one atomic orbital is usually denoted as a square cell:

Quantum mechanics has an extremely complex mathematical apparatus, therefore, in the framework of a school chemistry course, only the consequences of quantum mechanical theory are considered.

According to these consequences, any atomic orbital and the electron located in it are completely characterized by 4 quantum numbers.

• The principal quantum number, n, determines the total energy of an electron in a given orbital. The range of values ​​of the main quantum number is all natural numbers, i.e. n = 1,2,3,4, 5, etc.
• The orbital quantum number - l - characterizes the shape of the atomic orbital and can take any integer value from 0 to n-1, where n, recall, is the main quantum number.

Orbitals with l = 0 are called s-orbitals. s-Orbitals are spherical in shape and have no directionality in space:

Orbitals with l = 1 are called p-orbitals. These orbitals have the shape of a three-dimensional figure eight, i.e. a shape obtained by rotating a figure eight around an axis of symmetry, and outwardly resemble a dumbbell:

Orbitals with l = 2 are called d-orbitals, and with l = 3 – f-orbitals. Their structure is much more complex.

3) Magnetic quantum number – m l – determines the spatial orientation of a specific atomic orbital and expresses the projection of the orbital angular momentum onto the direction of the magnetic field. The magnetic quantum number m l corresponds to the orientation of the orbital relative to the direction of the external magnetic field strength vector and can take any integer values ​​from –l to +l, including 0, i.e. the total number of possible values ​​is (2l+1). So, for example, for l = 0 m l = 0 (one value), for l = 1 m l = -1, 0, +1 (three values), for l = 2 m l = -2, -1, 0, +1 , +2 (five values ​​of magnetic quantum number), etc.

So, for example, p-orbitals, i.e. orbitals with an orbital quantum number l = 1, having the shape of a “three-dimensional figure of eight,” correspond to three values ​​of the magnetic quantum number (-1, 0, +1), which, in turn, correspond to three directions perpendicular to each other in space.

4) The spin quantum number (or simply spin) - m s - can conditionally be considered responsible for the direction of rotation of the electron in the atom; it can take on values. Electrons with different spins are indicated by vertical arrows directed in different directions: ↓ and .

The set of all orbitals in an atom that have the same principal quantum number is called the energy level or electron shell. Any arbitrary energy level with some number n consists of n 2 orbitals.

A set of orbitals with the same values ​​of the principal quantum number and orbital quantum number represents an energy sublevel.

Each energy level, which corresponds to the principal quantum number n, contains n sublevels. In turn, each energy sublevel with orbital quantum number l consists of (2l+1) orbitals. Thus, the s sublevel consists of one s orbital, the p sublevel consists of three p orbitals, the d sublevel consists of five d orbitals, and the f sublevel consists of seven f orbitals. Since, as already mentioned, one atomic orbital is often denoted by one square cell, the s-, p-, d- and f-sublevels can be graphically represented as follows:

Each orbital corresponds to an individual strictly defined set of three quantum numbers n, l and m l.

The distribution of electrons among orbitals is called the electron configuration.

The filling of atomic orbitals with electrons occurs in accordance with three conditions:

• Minimum energy principle: Electrons fill orbitals starting from the lowest energy sublevel. The sequence of sublevels in increasing order of their energies is as follows: 1s<2s<2p<3s<3p<4s≤3d<4p<5s≤4d<5p<6s…;

To make it easier to remember this sequence of filling out electronic sublevels, the following graphic illustration is very convenient:

• Pauli principle: Each orbital can contain no more than two electrons.

If there is one electron in an orbital, then it is called unpaired, and if there are two, then they are called an electron pair.

• Hund's rule: the most stable state of an atom is one in which, within one sublevel, the atom has the maximum possible number of unpaired electrons. This most stable state of the atom is called the ground state.

In fact, the above means that, for example, the placement of 1st, 2nd, 3rd and 4th electrons in three orbitals of the p-sublevel will be carried out as follows:

The filling of atomic orbitals from hydrogen, which has a charge number of 1, to krypton (Kr), with a charge number of 36, will be carried out as follows:

Such a representation of the order of filling of atomic orbitals is called an energy diagram. Based on the electronic diagrams of individual elements, it is possible to write down their so-called electronic formulas (configurations). So, for example, an element with 15 protons and, as a consequence, 15 electrons, i.e. phosphorus (P) will have the following energy diagram:

When converted into an electronic formula, the phosphorus atom will take the form:

15 P = 1s 2 2s 2 2p 6 3s 2 3p 3

The normal size numbers to the left of the sublevel symbol show the energy level number, and the superscripts to the right of the sublevel symbol show the number of electrons in the corresponding sublevel.

Below are the electronic formulas of the first 36 elements of the periodic table by D.I. Mendeleev.

As already mentioned, in their ground state, electrons in atomic orbitals are located according to the principle of least energy. However, in the presence of empty p-orbitals in the ground state of the atom, often, by imparting excess energy to it, the atom can be transferred to the so-called excited state. For example, a boron atom in its ground state has an electronic configuration and an energy diagram of the following form:

And in an excited state (*), i.e. When some energy is imparted to a boron atom, its electron configuration and energy diagram will look like this:

Depending on which sublevel in the atom is filled last, chemical elements are divided into s, p, d or f.

Finding s, p, d and f elements in the table D.I. Mendeleev:

• The s-elements have the last s-sublevel to be filled. These elements include elements of the main (on the left in the table cell) subgroups of groups I and II.
• For p-elements, the p-sublevel is filled. The p-elements include the last six elements of each period, except the first and seventh, as well as elements of the main subgroups of groups III-VIII.
• d-elements are located between s- and p-elements in large periods.
• f-Elements are called lanthanides and actinides. They are listed at the bottom of the D.I. table. Mendeleev.

Since during chemical reactions the nuclei of the reacting atoms remain unchanged (with the exception of radioactive transformations), the chemical properties of atoms depend on the structure of their electronic shells. Theory electronic structure of the atom built on the basis of the apparatus of quantum mechanics. Thus, the structure of atomic energy levels can be obtained on the basis of quantum mechanical calculations of the probabilities of finding electrons in the space around the atomic nucleus ( rice. 4.5).

Rice. 4.5. Scheme of dividing energy levels into sublevels

The fundamentals of the theory of the electronic structure of an atom are reduced to the following provisions: the state of each electron in an atom is characterized by four quantum numbers: the principal quantum number n = 1, 2, 3,; orbital (azimuthal) l=0,1,2, n–1; magnetic m l = –l, –1,0,1,  l; spin m s = -1/2, 1/2 .

According to Pauli principle, in the same atom there cannot be two electrons having the same set of four quantum numbers n, l, m l , m s; collections of electrons with the same principal quantum numbers n form electron layers, or energy levels of the atom, numbered from the nucleus and denoted as K, L, M, N, O, P, Q, and in the energy layer with a given value n can be no more than 2n 2 electrons. Collections of electrons with the same quantum numbers n And l, form sublevels, designated as they move away from the core as s, p, d, f.

The probabilistic determination of the position of the electron in space around the atomic nucleus corresponds to the Heisenberg uncertainty principle. According to quantum mechanical concepts, an electron in an atom does not have a specific trajectory of motion and can be located in any part of the space around the nucleus, and its various positions are considered as an electron cloud with a certain negative charge density. The space around the nucleus in which an electron is most likely to be found is called orbital. It contains about 90% of the electron cloud. Each sublevel 1s, 2s, 2p etc. corresponds to a certain number of orbitals of a certain shape. For example, 1s- And 2s- orbitals are spherical and 2p-orbitals ( 2p x , 2p y , 2p z-orbitals) are oriented in mutually perpendicular directions and have the shape of a dumbbell ( rice. 4.6).

Rice. 4.6. Shape and orientation of electron orbitals.

During chemical reactions, the atomic nucleus does not undergo changes; only the electronic shells of the atoms change, the structure of which explains many of the properties of chemical elements. Based on the theory of the electronic structure of the atom, the deep physical meaning of Mendeleev’s periodic law of chemical elements was established and the theory of chemical bonding was created.

The theoretical justification of the periodic system of chemical elements includes data on the structure of the atom, confirming the existence of a connection between the periodicity of changes in the properties of chemical elements and the periodic repetition of similar types of electronic configurations of their atoms.

In the light of the doctrine of the structure of the atom, Mendeleev’s division of all elements into seven periods becomes justified: the number of the period corresponds to the number of energy levels of atoms filled with electrons. In small periods, with an increase in the positive charge of atomic nuclei, the number of electrons at the external level increases (from 1 to 2 in the first period, and from 1 to 8 in the second and third periods), which explains the change in the properties of elements: at the beginning of the period (except for the first) there is alkali metal, then a gradual weakening of metallic properties and strengthening of non-metallic properties is observed. This pattern can be traced for elements of the second period in table 4.2.

Table 4.2.

In large periods, as the charge of the nuclei increases, the filling of levels with electrons is more difficult, which explains the more complex change in the properties of elements compared to elements of small periods.

The identical nature of the properties of chemical elements in subgroups is explained by the similar structure of the external energy level, as shown in table 4.3, illustrating the sequence of filling energy levels with electrons for subgroups of alkali metals.

Table 4.3.

The group number usually indicates the number of electrons in an atom that can participate in the formation of chemical bonds. This is the physical meaning of the group number. In four places of the periodic table, the elements are not arranged in order of increasing atomic mass: Ar And K,Co And Ni,Te And I,Th And Pa. These deviations were considered shortcomings of the periodic table of chemical elements. The doctrine of the structure of the atom explained these deviations. Experimental determination of nuclear charges showed that the arrangement of these elements corresponds to an increase in the charges of their nuclei. In addition, the experimental determination of the charges of atomic nuclei made it possible to determine the number of elements between hydrogen and uranium, as well as the number of lanthanides. Now all places in the periodic table are filled in the interval from Z=1 before Z=114, however, the periodic system is not complete, the discovery of new transuranium elements is possible.

Electrons

The concept of atom arose in the ancient world to designate particles of matter. Translated from Greek, atom means “indivisible.”

The Irish physicist Stoney, based on experiments, came to the conclusion that electricity is carried by the smallest particles existing in the atoms of all chemical elements. In 1891, Stoney proposed to call these particles electrons, which means “amber” in Greek. A few years after the electron got its name, the English physicist Joseph Thomson and the French physicist Jean Perrin proved that electrons carry a negative charge. This is the smallest negative charge, which in chemistry is taken as one (-1). Thomson even managed to determine the speed of the electron (the speed of the electron in the orbit is inversely proportional to the orbit number n. The radii of the orbits increase in proportion to the square of the orbit number. In the first orbit of the hydrogen atom (n=1; Z=1) the speed is ≈ 2.2·106 m/ s, that is, about a hundred times less than the speed of light c = 3·108 m/s) and the mass of the electron (it is almost 2000 times less than the mass of the hydrogen atom).

State of electrons in an atom

The state of an electron in an atom is understood as a set of information about the energy of a particular electron and the space in which it is located. An electron in an atom does not have a trajectory of motion, i.e. we can only talk about the probability of finding it in the space around the nucleus.

It can be located in any part of this space surrounding the nucleus, and the totality of its various positions is considered as an electron cloud with a certain negative charge density. Figuratively, this can be imagined this way: if it were possible to photograph the position of an electron in an atom after hundredths or millionths of a second, as in a photo finish, then the electron in such photographs would be represented as dots. If countless such photographs were superimposed, the picture would be of an electron cloud with the greatest density where there would be the most of these points.

The space around the atomic nucleus in which an electron is most likely to be found is called an orbital. It contains approximately 90% electronic cloud, and this means that about 90% of the time the electron is in this part of space. They are distinguished by shape 4 currently known types of orbitals, which are designated by Latin letters s, p, d and f. A graphical representation of some forms of electron orbitals is presented in the figure.

The most important characteristic of the motion of an electron in a certain orbital is energy of its connection with the nucleus. Electrons with similar energy values ​​form a single electron layer, or energy level. Energy levels are numbered starting from the nucleus - 1, 2, 3, 4, 5, 6 and 7.

The integer n, indicating the number of the energy level, is called the principal quantum number. It characterizes the energy of electrons occupying a given energy level. Electrons of the first energy level, closest to the nucleus, have the lowest energy. Compared to electrons of the first level, electrons of subsequent levels will be characterized by a large supply of energy. Consequently, the electrons of the outer level are least tightly bound to the atomic nucleus.

The largest number of electrons at an energy level is determined by the formula:

N = 2n 2 ,

where N is the maximum number of electrons; n is the level number, or the main quantum number. Consequently, at the first energy level closest to the nucleus there can be no more than two electrons; on the second - no more than 8; on the third - no more than 18; on the fourth - no more than 32.

Starting from the second energy level (n = 2), each of the levels is divided into sublevels (sublayers), slightly different from each other in the binding energy with the nucleus. The number of sublevels is equal to the value of the main quantum number: the first energy level has one sublevel; the second - two; third - three; fourth - four sublevels. The sublevels, in turn, are formed by orbitals. Each valuen corresponds to the number of orbitals equal to n.

Sublevels are usually denoted by Latin letters, as well as the shape of the orbitals of which they consist: s, p, d, f.

Protons and Neutrons

An atom of any chemical element is comparable to a tiny solar system. Therefore, this model of the atom, proposed by E. Rutherford, is called planetary.

The atomic nucleus, in which the entire mass of the atom is concentrated, consists of particles of two types - protons and neutrons.

Protons have a charge equal to the charge of electrons, but opposite in sign (+1), and a mass equal to the mass of a hydrogen atom (it is taken as one in chemistry). Neutrons carry no charge, they are neutral and have a mass equal to the mass of a proton.

Protons and neutrons together are called nucleons (from the Latin nucleus - nucleus). The sum of the number of protons and neutrons in an atom is called the mass number. For example, the mass number of an aluminum atom is:

13 + 14 = 27

number of protons 13, number of neutrons 14, mass number 27

Since the mass of the electron, which is negligibly small, can be neglected, it is obvious that the entire mass of the atom is concentrated in the nucleus. Electrons are designated e - .

Since the atom electrically neutral, then it is also obvious that the number of protons and electrons in an atom is the same. It is equal to the serial number of the chemical element assigned to it in the Periodic Table. The mass of an atom consists of the mass of protons and neutrons. Knowing the atomic number of the element (Z), i.e. the number of protons, and the mass number (A), equal to the sum of the numbers of protons and neutrons, you can find the number of neutrons (N) using the formula:

N = A - Z

For example, the number of neutrons in an iron atom is:

56 — 26 = 30

Isotopes

Varieties of atoms of the same element that have the same nuclear charge but different mass numbers are called isotopes. Chemical elements found in nature are a mixture of isotopes. Thus, carbon has three isotopes with masses 12, 13, 14; oxygen - three isotopes with masses 16, 17, 18, etc. The relative atomic mass of a chemical element usually given in the Periodic Table is the average value of the atomic masses of a natural mixture of isotopes of a given element, taking into account their relative abundance in nature. The chemical properties of isotopes of most chemical elements are exactly the same. However, hydrogen isotopes vary greatly in properties due to the dramatic multiple increase in their relative atomic mass; they are even given individual names and chemical symbols.

Elements of the first period

Diagram of the electronic structure of the hydrogen atom:

Diagrams of the electronic structure of atoms show the distribution of electrons across electronic layers (energy levels).

Graphic electronic formula of the hydrogen atom (shows the distribution of electrons by energy levels and sublevels):

Graphic electronic formulas of atoms show the distribution of electrons not only among levels and sublevels, but also among orbitals.

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.

For all elements of the second period the first electronic layer is filled, and electrons fill the s- 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.

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

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.

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

In aluminum and subsequent elements, the 3p sublevel is filled with electrons.

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.

Elements of the fourth - seventh periods

A fourth electron layer appears in potassium and calcium atoms, and the 4s sublevel is filled, since it has lower energy than the 3d sublevel.

K, Ca - s-elements included in the main subgroups. For atoms from Sc to Zn, the 3d sublevel is filled with electrons. These are 3d elements. They are included in 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, one electron “fails” from the 4s to the 3d sublevel, which is explained by the greater energy stability of the resulting electronic configurations 3d 5 and 3d 10:

In the zinc atom, the third electron layer is complete - all 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 there can be a total of 32 electrons in the fourth electron layer; the krypton atom still has unfilled 4d and 4f sublevels. For elements of the fifth period, sublevels are being filled in the following order: 5s - 4d - 5p. And there are also exceptions related to “ failure» electrons, y 41 Nb, 42 Mo, 44 ​​Ru, 45 Rh, 46 Pd, 47 Ag.

In the sixth and seventh periods, f-elements appear, i.e., 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 the atoms of elements of the sixth period: 55 Cs and 56 Ba - 6s elements; 57 La … 6s 2 5d x - 5d element; 58 Ce - 71 Lu - 4f elements; 72 Hf - 80 Hg - 5d elements; 81 T1 - 86 Rn - 6d elements. But here, too, there are elements in which the order of filling the electronic orbitals is “violated,” which, for example, is associated with the greater energy stability of half and fully filled f-sublevels, i.e. nf 7 and nf 14. Depending on which sublevel of the atom is filled with electrons last, all elements are divided into four electron families, or blocks:

• s-elements. The s-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.

• 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.

• 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, i.e. elements of plug-in decades of large periods located between s- and p-elements. They are also called transition elements.

• f-elements. The f-sublevel of the third outer level of the atom is filled with electrons; these include lanthanides and antinoids.

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”), i.e., having such properties that conditionally can be imagined as the rotation of an electron around its imaginary axis: clockwise or counterclockwise.

This principle is called Pauli principle. If there is one electron in the orbital, then it is called unpaired; if there are two, then these are paired electrons, i.e. electrons with opposite spins. The figure shows a diagram of the division of energy levels into sublevels and the order in which they are filled.

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: Pauli's principle and F. Hund's rule, according to which electrons occupy free cells first one at a time and have the same spin value, and only then pair, but the spins, according to the Pauli principle, will already be oppositely directed.

Hund's rule and Pauli's principle

Hund's rule- a rule of quantum chemistry that determines the order of filling the orbitals of a certain sublayer and is formulated as follows: the total value of the spin quantum number of electrons of a given sublayer must be maximum. Formulated by Friedrich Hund in 1925.

This means that in each of the orbitals of the sublayer, one electron is filled first, and only after the unfilled orbitals are exhausted, a second electron is added to this orbital. In this case, in one orbital there are two electrons with half-integer spins of the opposite sign, which pair (form a two-electron cloud) and, as a result, the total spin of the orbital becomes equal to zero.

Another wording: Lower in energy lies the atomic term for which two conditions are satisfied.

1. Multiplicity is maximum
2. When the multiplicities coincide, the total orbital momentum L is maximum.

Let us analyze this rule using the example of filling p-sublevel orbitals p-elements of the second period (that is, from boron to neon (in the diagram below, horizontal lines indicate orbitals, vertical arrows indicate electrons, and the direction of the arrow indicates the spin orientation).

Klechkovsky's rule

Klechkovsky's rule - as the total number of electrons in atoms increases (with an increase in the charges of their nuclei, or the serial numbers of chemical elements), atomic orbitals are populated in such a way that the appearance of electrons in an orbital with a higher energy depends only on the main quantum number n and does not depend on all other quantum numbers numbers, including from l. Physically, this means that in a hydrogen-like atom (in the absence of interelectron repulsion), the orbital energy of an electron is determined only by the spatial distance of the electron charge density from the nucleus and does not depend on the characteristics of its motion in the field of the nucleus.

The empirical Klechkovsky rule and the ordering scheme that follows from it are somewhat contradictory to the real energy sequence of atomic orbitals only in two similar cases: for atoms Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, Pt, Au, there is a “failure” of an electron with s -sublevel of the outer layer is replaced by the d-sublevel of the previous layer, which leads to an energetically more stable state of the atom, namely: after filling orbital 6 with two electrons s

An atom is the smallest particle of matter, consisting of a nucleus and electrons. The structure of the electronic shells of atoms is determined by the position of the element in the Periodic Table of Chemical Elements by D.I. Mendeleev.

Electron and electron shell of an atom

An atom, which is generally neutral, consists of a positively charged nucleus and a negatively charged electron shell (electron cloud), with the total positive and negative charges being equal in absolute value. When calculating the relative atomic mass, the mass of electrons is not taken into account, since it is negligible and 1840 times less than the mass of a proton or neutron.

Rice. 1. Atom.

An electron is a completely unique particle that has a dual nature: it has both the properties of a wave and a particle. They continuously move around the core.

The space around the nucleus where the probability of finding an electron is most likely is called an electron orbital, or electron cloud. This space has a specific shape, which is designated by the letters s-, p-, d-, and f-. The S-electron orbital has a spherical shape, the p-orbital has the shape of a dumbbell or a three-dimensional figure eight, the shapes of the d- and f-orbitals are much more complex.

Rice. 2. Shapes of electron orbitals.

Around the nucleus, electrons are arranged in electron layers. Each layer is characterized by its distance from the nucleus and its energy, which is why electronic layers are often called electronic energy levels. The closer the level is to the nucleus, the lower the energy of the electrons in it. One element differs from another in the number of protons in the nucleus of the atom and, accordingly, in the number of electrons. Consequently, the number of electrons in the electron shell of a neutral atom is equal to the number of protons contained in the nucleus of this atom. Each subsequent element has one more proton in its nucleus, and one more electron in its electron shell.

The newly entering electron occupies the orbital with the lowest energy. However, the maximum number of electrons per level is determined by the formula:

where N is the maximum number of electrons, and n is the number of the energy level.

The first level can only have 2 electrons, the second can have 8 electrons, the third can have 18 electrons, and the fourth level can have 32 electrons. The outer level of an atom cannot contain more than 8 electrons: as soon as the number of electrons reaches 8, the next level, further from the nucleus, begins to be filled.

Structure of electronic shells of atoms

Each element stands in a certain period. A period is a horizontal collection of elements arranged in order of increasing charge of the nuclei of their atoms, which begins with an alkali metal and ends with an inert gas. The first three periods in the table are small, and the next, starting from the fourth period, are large, consisting of two rows. The number of the period in which the element is located has a physical meaning. It means how many electronic energy levels there are in an atom of any element of a given period. Thus, the element chlorine Cl is in the 3rd period, that is, its electron shell has three electronic layers. Chlorine is in group VII of the table, and in the main subgroup. The main subgroup is the column within each group that begins with period 1 or 2.

Thus, the state of the electron shells of the chlorine atom is as follows: the atomic number of the chlorine element is 17, which means that the atom has 17 protons in the nucleus and 17 electrons in the electron shell. At level 1 there can only be 2 electrons, at level 3 - 7 electrons, since chlorine is in the main subgroup of group VII. Then at level 2 there are: 17-2-7 = 8 electrons.