The structure of the electronic levels of an atom. Atomic structure, chemical bonding, valency and molecular structure

(Lecture notes)

The structure of the atom. Introduction.

The object of study in chemistry is chemical elements and their compounds. Chemical element called a collection of atoms with the same positive charge. Atom- This smallest particle chemical element, preserving it Chemical properties. Bonding with each other, atoms of one or different elements form more complex particles - molecules. A collection of atoms or molecules form chemical substances. Each individual chemical substance is characterized by a set of individual physical properties, such as boiling and melting points, density, electrical and thermal conductivity, etc.

1. Atomic structure and the Periodic Table of Elements

DI. Mendeleev.

Knowledge and understanding of the laws of the order of filling the Periodic Table of Elements D.I. Mendeleev allows us to understand the following:

1. the physical essence of the existence of certain elements in nature,

2. the nature of the chemical valence of the element,

3. the ability and “lightness” of an element to give or accept electrons when interacting with another element,

4. the nature of the chemical bonds that a given element can form when interacting with other elements, the spatial structure of simple and complex molecules, etc., etc.

The structure of the atom.

An atom is a complex microsystem of things in motion and interacting with each other elementary particles.

In the late 19th and early 20th centuries, it was discovered that atoms are made up of smaller particles: neutrons, protons and electrons. The last two particles are charged particles, the proton carries a positive charge, the electron a negative charge. Since the atoms of an element in the ground state are electrically neutral, this means that the number of protons in an atom of any element is equal to the number of electrons. The mass of atoms is determined by the sum of the masses of protons and neutrons, the number of which is equal to the difference between the mass of the atoms and its serial number in periodic table DI. Mendeleev.

In 1926, Schrödinger proposed describing the movement of microparticles in the atom of an element using the wave equation he derived. When solving the Schrödinger wave equation for the hydrogen atom, three integer quantum numbers appear: n, ℓ And m ℓ , which characterize the state of the electron in three-dimensional space in the central field of the nucleus. Quantum numbers n, ℓ And m ℓ take integer values. Wave function defined by three quantum numbers n, ℓ And m ℓ and obtained as a result of solving the Schrödinger equation is called an orbital. An orbital is a region of space in which an electron is most likely to be found, belonging to an atom of a chemical element. Thus, solving the Schrödinger equation for the hydrogen atom leads to the appearance of three quantum numbers, physical meaning which is that they characterize three different types of orbitals that an atom can have. Let's take a closer look at each quantum number.

Principal quantum number n can take any positive integer values: n = 1,2,3,4,5,6,7...It characterizes the energy of the electron level and the size of the electron “cloud”. It is characteristic that the number of the main quantum number coincides with the number of the period in which the element is located.

Azimuthal or orbital quantum numberℓ can take integer values ​​from ℓ = 0….to n – 1 and determines the moment of electron motion, i.e. orbital shape. For various numerical valuesℓ use the following notation: ℓ = 0, 1, 2, 3, and are indicated by the symbols s, p, d, f, respectively for ℓ = 0, 1, 2 and 3. In the periodic table of elements there are no elements with a spin number ℓ = 4.

Magnetic quantum numberm ℓ characterizes the spatial arrangement of electron orbitals and, consequently, the electromagnetic properties of the electron. It can take values ​​from – ℓ to + ℓ , including zero.

The shape, or more precisely, the symmetry properties of atomic orbitals depend on quantum numbers ℓ And m ℓ . "Electronic cloud" corresponding s- the orbitals have, have the shape of a ball (at the same time ℓ = 0).

Fig.1. 1s orbital

The orbitals defined by the quantum numbers ℓ = 1 and m ℓ = -1, 0 and +1 are called p-orbitals. Since m ℓ in this case has three different meanings, then the atom has three energetically equivalent p-orbitals (the principal quantum number for them is the same and can have the value n = 2,3,4,5,6 or 7). p-Orbitals have axial symmetry and look like three-dimensional figure eights, oriented along the x, y and z axes in an external field (Fig. 1.2). Hence the origin of the symbolism p x , p y and p z .

Fig.2. p x, p y and p z orbitals

In addition, there are d- and f- atomic orbitals, for the first ℓ = 2 and m ℓ = -2, -1, 0, +1 and +2, i.e. five AOs, for the second ones ℓ = 3 and m ℓ = -3, -2, -1, 0, +1, +2 and +3, i.e. 7 JSC.

Fourth quantum m s called the spin quantum number, was introduced to explain certain subtle effects in the spectrum of the hydrogen atom by Goudsmit and Uhlenbeck in 1925. The spin of an electron is the angular momentum of a charged elementary particle of an electron, the orientation of which is quantized, i.e. strictly limited to certain angles. This orientation is determined by the value of the spin magnetic quantum number (s), which for the electron is equal to ½ , therefore for the electron according to the quantization rules m s = ± ½. In this regard, to the set of three quantum numbers we should add the quantum number m s . Let us emphasize once again that four quantum numbers determine the order of construction of Mendeleev’s periodic table of elements and explain why there are only two elements in the first period, eight in the second and third, 18 in the fourth, etc. However, in order to explain the structure of many-electron atoms, the order of filling electronic levels as the positive charge of the atom increases, it is not enough to have an idea of ​​the four quantum numbers that “control” the behavior of electrons when filling electron orbitals, but you need to know some more simple rules, namely, Pauli's principle, Hund's rule and Kleczkowski's rules.

According to the Pauli principle In the same quantum state, characterized by certain values ​​of four quantum numbers, there cannot be more than one electron. This means that one electron can, in principle, be placed in any atomic orbital. Two electrons can be in the same atomic orbital only if they have different spin quantum numbers.

When filling three p-AOs, five d-AOs and seven f-AOs with electrons, one should be guided, in addition to the Pauli principle, by Hund’s rule: The filling of the orbitals of one subshell in the ground state occurs with electrons with identical spins.

When filling the subshells (p, d, f)the absolute value of the sum of spins must be maximum.

Klechkovsky's rule. According to Klechkovsky’s rule, when fillingd And felectron orbital must be respectedprinciple of minimum energy. According to this principle, electrons in the ground state fill orbitals with minimum levels energy. The energy of a sublevel is determined by the sum of quantum numbersn + ℓ = E .

Klechkovsky's first rule: First, those sublevels for whichn + ℓ = E minimal.

Klechkovsky's second rule: in case of equalityn + ℓ for several sublevels, the sublevel for which is filledn minimal .

Currently, 109 elements are known.

2. Ionization energy, electron affinity and electronegativity.

The most important characteristics of the electronic configuration of an atom are ionization energy (IE) or ionization potential (IP) and the atom's electron affinity (EA). Ionization energy is the change in energy during the removal of an electron from a free atom at 0 K: A = + + ē . The dependence of ionization energy on the atomic number Z of an element and the size of the atomic radius has a pronounced periodic character.

Electron affinity (EA) is the change in energy that accompanies the addition of an electron to an isolated atom to form a negative ion at 0 K: A + ē = A - (the atom and ion are in their ground states). In this case, the electron occupies the lowest vacant atomic orbital (LUAO) if the VZAO is occupied by two electrons. The SE strongly depends on their orbital electronic configuration.

Changes in EI and SE correlate with changes in many properties of elements and their compounds, which is used to predict these properties from EI and SE values. The highest absolute value Halogens have an affinity for electrons. In each group periodic table elements, the ionization potential or EI decreases with increasing element number, which is associated with an increase in atomic radius and with an increase in the number of electronic layers and which correlates well with an increase in the reducing power of the element.

Table 1 of the Periodic Table of Elements shows the values ​​of EI and SE in eV/per atom. Note that exact values SEs are known only for a few atoms; their values ​​are highlighted in Table 1.

Table 1

First ionization energy (EI), electron affinity (EA) and electronegativity χ) of atoms in the periodic table.

χ	0.747 2. 1 0 0, 3 7																	1,2 2
χ	0.54 1. 55	-0.3 1. 1 3											0.2 0. 91	1.2 5	-0. 1 0, 55	1.47 0. 59	3.45 0. 64	1 ,60
χ	0. 7 4 1. 89	-0.3 1 . 3 1 1 . 6 0											0. 6	1.63	0.7	2.07	3.61
χ	2.3 6	- 0 .6						1.26(α)				-0.9 1 . 39	0. 18	1.2	0. 6	2.07	3.36
χ	2.4 8											-0.6 1 . 56	0. 2			2.2
χ	2.6 7	2, 2 1						ABOUTs

χ – electronegativity according to Pauling

r- atomic radius, (from “Laboratory and seminar classes in general and inorganic chemistry”, N.S. Akhmetov, M.K. Azizova, L.I. Badygina)

Electrons

The concept of atom arose in ancient world to designate particles of matter. Translated from Greek atom means "indivisible".

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, English physicist Joseph Thomson and 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 density negative charge. 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.

Space around atomic nucleus, in which the 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. Graphic image Some forms of electron orbitals are shown 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, electrons are least tightly bound to the nucleus of an atom external level.

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

N = 2n 2 ,

where N - maximum number 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 designated with Latin letters, as well as the shape of the orbitals of which they are composed: 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 charge electrons, but opposite in sign (+1), and mass, equal to mass hydrogen atom (it is taken as a unit 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 serial number 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 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 same charge nuclei 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 this 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 electronic shells 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 filled, respectively.

4f elements are called lanthanides.

5f elements are called actinides.

Filling procedure electronic sublevels in 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 at the same time have same value back, and only then mate, but the backs, according to the Pauli principle, will already be in opposite directions.

Hund's rule and Pauli's principle

Hund's rule- rule quantum chemistry, which determines the order of filling the orbitals of a certain sublayer and is formulated in the following way: total value the spin quantum number of electrons of a given sublayer should 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, one orbital contains two electrons with half-integer spins 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 orbitals with more high energy depends only on the principal quantum number n and does not depend on all other quantum numbers, including 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 to the d-sublevel of the previous layer, which leads to an energetically more steady state atom, namely: after filling orbital 6 with two electrons s

The composition of the molecule. That is, what atoms form the molecule, in what quantity, and by what bonds these atoms are connected. All this determines the property of the molecule, and accordingly the property of the substance that these molecules form.

For example, the properties of water: transparency, fluidity, and the ability to cause rust are due precisely to the presence of two hydrogen atoms and one oxygen atom.

Therefore, before we begin to study the properties of molecules (that is, the properties of substances), we need to consider the “building blocks” with which these molecules are formed. Understand the structure of the atom.

How is an atom structured?

Atoms are particles that combine with each other to form molecules.

The atom itself consists of positively charged nucleus (+) And negatively charged electron shell (-). In general, the atom is electrically neutral. That is, the charge of the nucleus is equal in absolute value to the charge of the electron shell.

The nucleus is formed by the following particles:

• Protons. One proton carries a +1 charge. Its mass is 1 amu (atomic mass unit). These particles are necessarily present in the nucleus.

• Neutrons. The neutron has no charge (charge = 0). Its mass is 1 amu. There may be no neutrons in the nucleus. It is not an essential component of the atomic nucleus.

Thus, protons are responsible for the overall charge of the nucleus. Since one neutron has a charge of +1, the charge of the nucleus is equal to the number of protons.

The electron shell, as the name suggests, is formed by particles called electrons. If we compare the nucleus of an atom with a planet, then electrons are its satellites. Rotating around the nucleus (for now let’s imagine that in orbits, but in fact in orbitals), they form an electron shell.

• Electron- This is a very small particle. Its mass is so small that it is taken as 0. But the charge of the electron is -1. That is, modulo equal to charge proton, differs in sign. Since one electron carries a -1 charge, the total charge of the electron shell is equal to the number of electrons in it.

One important consequence is that since an atom is a particle that has no charge (the charge of the nucleus and the charge of the electron shell are equal in magnitude, but opposite in sign), that is, electrically neutral, therefore, the number of electrons in an atom is equal to the number of protons.

How do atoms of different chemical elements differ from each other?

Atoms of different chemical elements differ from each other in the charge of the nucleus (that is, the number of protons, and, consequently, the number of electrons).

How to find out the charge of the nucleus of an atom of an element? The brilliant domestic chemist D.I. Mendeleev, having discovered periodic law, and by developing a table named after him, gave us the opportunity to do this. His discovery was far ahead. When the structure of the atom was not yet known, Mendeleev arranged the elements in the table in order of increasing nuclear charge.

That is, the serial number of an element in the periodic table is the charge of the nucleus of an atom of a given element. For example, oxygen has a serial number of 8, so the charge on the nucleus of an oxygen atom is +8. Accordingly, the number of protons is 8, and the number of electrons is 8.

It is the electrons in the electron shell that determine Chemical properties atom, but more on that a little later.

Now let's talk about mass.

One proton is one unit of mass, one neutron is also one unit of mass. Therefore, the sum of neutrons and protons in a nucleus is called mass number. (Electrons do not affect the mass in any way, since we neglect its mass and consider it equal to zero).

Atomic unit mass (a.m.u.) – special physical quantity to denote the small masses of particles that form atoms.

All these three atoms are atoms of one chemical element - hydrogen. Because they have the same nuclear charge.

How will they be different? These atoms have different mass numbers (due to different numbers neutrons). The first atom has a mass number of 1, the second has 2, and the third has 3.

Atoms of the same element that differ in the number of neutrons (and therefore mass numbers) are called isotopes.

The presented hydrogen isotopes even have their own names:

• The first isotope (with mass number 1) is called protium.
• The second isotope (with mass number 2) is called deuterium.
• The third isotope (with mass number 3) is called tritium.

Now the next reasonable question: why if the number of neutrons and protons in the nucleus is an integer, their mass is 1 amu, then in the periodic system the mass of an atom is a fractional number. For sulfur, for example: 32.066.

Answer: the element has several isotopes, they differ from each other in mass numbers. Therefore, the atomic mass in the periodic table is the average value of the atomic masses of all isotopes of an element, taking into account their occurrence in nature. This mass, indicated in the periodic table, is called relative atomic mass.

For chemical calculations, the indicators of just such an “average atom” are used. Atomic mass rounded to the nearest whole number.

The structure of the electron shell.

The chemical properties of an atom are determined by the structure of its electron shell. Electrons around the nucleus are not located anyhow. Electrons are localized in electron orbitals.

Electron orbital– the space around the atomic nucleus where the probability of finding an electron is greatest.

An electron has one quantum parameter called spin. If you take classic definition from quantum mechanics, That spin- This own moment momentum of the particle. In a simplified form, this can be represented as the direction of rotation of a particle around its axis.

An electron is a particle with half-integer spin; an electron can have either +½ or -½ spin. Conventionally, this can be represented as clockwise and counterclockwise rotation.

One electron orbital can contain no more than two electrons with opposite spins.

The generally accepted designation for an electronic habitat is a cell or a dash. An electron is designated by an arrow: an up arrow is an electron with a positive spin +½, a down arrow ↓ is an electron with a negative spin -½.

An electron alone in an orbital is called unpaired. Two electrons located in the same orbital are called paired.

Electronic orbitals are divided into four types depending on their shape: s, p, d, f. Orbitals same shape form a sublevel. The number of orbitals at a sublevel is determined by the number possible options location in space.

1. s-orbital.

The s-orbital has the shape of a ball:

In space, the s-orbital can be located in only one way:

Therefore, the s sublevel is formed by only one s orbital.

1. p-orbital.

The p-orbital is shaped like a dumbbell:

In space, the p-orbital can be located in only three ways:

Therefore, the p-sublevel is formed by three p-orbitals.

1. d-orbital.

d-orbital has complex shape:

In space, the d-orbital can be arranged in five different ways. Therefore, the d sublevel is formed by five d orbitals.

1. f-orbital

The f orbital has an even more complex shape. In space, the f orbital can be located in seven different ways. Therefore, the f sublevel is formed by seven f orbitals.

The electron shell of an atom is like a puff pastry product. It also has layers. Electrons located on different layers have different energies: on layers closer to the nucleus they have less energy, on layers farther from the nucleus they have more energy. These layers are called energy levels.

Filling electron orbitals.

The first energy level has only the s-sublevel:

At the second energy level there is an s-sublevel and a p-sublevel appears:

At the third energy level there is an s-sublevel, a p-sublevel, and a d-sublevel appears:

At the fourth energy level, in principle, an f-sublevel is added. But in school course The f orbitals are not filled, so we don't have to draw the f sublevel:

The number of energy levels in an atom of an element is period number. When filling electron orbitals, you must follow the following principles:

1. Each electron tries to occupy the position in the atom where its energy is minimal. That is, first the first energy level is filled, then the second, and so on.

The electronic formula is also used to describe the structure of the electron shell. Electronic formula is a short one-line summary of the distribution of electrons across sublevels.

1. At a sublevel, each electron first fills an empty orbital. And each has spin +½ (up arrow).

And only after each sublevel orbital has one electron, the next electron becomes paired - that is, it occupies an orbital that already has an electron:

1. The d-sublevel is filled in a special way.

The fact is that the energy of the d-sublevel is higher than the energy of the s-sublevel of the NEXT energy layer. And as we know, the electron tries to occupy that position in the atom where its energy will be minimal.

Therefore, after filling the 3p sublevel, the 4s sublevel is filled first, after which the 3d sublevel is filled.

And only after the 3d sublevel is completely filled, the 4p sublevel is filled.

The same goes for energy level 4. After filling the 4p sublevel, the 5s sublevel is filled next, followed by the 4d sublevel. And after it only 5p.

1. And there is one more point, one rule regarding filling out the d-sublevel.

Then a phenomenon occurs called failure. If there is a failure, one electron from the s-sublevel of the next energy level, in literally falls to the d-electron.

Ground and excited states of the atom.

atoms, electronic configurations which we have now built are called atoms in basic condition. That is, this is a normal, natural, if you like, state.

When an atom receives energy from outside, excitation can occur.

Excitation is the transition of a paired electron to an empty orbital, within the outer energy level.

For example, for a carbon atom:

Excitation is characteristic of many atoms. This must be remembered because excitation determines the ability of atoms to bond with each other. The main thing to remember is the condition under which excitation can occur: a paired electron and an empty orbital at the outer energy level.

There are atoms that have several excited states:

Electronic configuration of the ion.

Ions are particles into which atoms and molecules turn by gaining or losing electrons. These particles have a charge because they either have a “lack” of electrons or an excess of them. Positively charged ions are called cations, negative – anions.

The chlorine atom (has no charge) gains an electron. An electron has a charge of 1- (one minus), and accordingly a particle is formed that has an excess negative charge. Chlorine anion:

Cl 0 + 1e → Cl –

The lithium atom (also having no charge) loses an electron. The electron has a charge of 1+ (one plus), a particle is formed with a lack of negative charge, that is, it has a positive charge. Lithium cation:

Li 0 – 1e → Li +

Transforming into ions, atoms acquire such a configuration that the outer energy level becomes “beautiful,” that is, completely filled. This configuration is the most thermodynamically stable, so there is a reason for atoms to turn into ions.

And therefore the atoms of the elements VIII-A group(eighth group main subgroup), as stated in the next paragraph, these are noble gases and are chemically inactive. Their basic state has the following structure: the outer energy level is completely filled. Other atoms seem to strive to acquire the configuration of these most noble gases, and therefore turn into ions and form chemical bonds.

Mendeleev's periodic table of elements. The structure of the atom.

MENDELEEV'S PERIODIC SYSTEM OF ELEMENTS - chemical classification. elements created by Russian. scientist D.I. Mendeleev on the basis of the periodicity discovered by him (in 1869). law.

Modern periodic formulation law: the properties of elements (manifested in simple compounds and compounds) are found in periodic periods. depending on the charge of the nuclei of their atoms.

The charge of the atomic nucleus Z is equal to the atomic (ordinal) number of the chemical. element in P. s. e. M. If you arrange all the elements in ascending order Z. (hydrogen H, Z = 1; helium He, Z = 2; lithium Li, Z == 3; beryllium Be, Z = 4, etc.), then they form 7 periods. In each of these periods, a regular change in the properties of the elements is observed, from the first element of the period (alkali metal) to the last (noble gas). The first period contains 2 elements, the 2nd and 3rd - 8 elements each, the 4th and 5th - 18, the 6th - 32. In the 7th period, 19 elements are known. The 2nd and 3rd periods are usually called small, all subsequent periods are called large. If you arrange the periods in the form of horizontal rows, then the resulting the table will show 8 vertical lines. columns; These are groups of elements that are similar in their properties.

The properties of elements within groups also naturally change depending on the increase in Z. For example, in the group Li - Na - K - Rb - Cs - Fr, the chemical content increases. the activity of the metal is enhanced by nature of oxides and hydroxides.

From the theory of atomic structure it follows that the periodicity of the properties of elements is determined by the laws of the formation of electron shells around the nucleus. As the element's Z increases, the atom becomes more complex - the number of electrons surrounding the nucleus increases, and a moment comes when the filling of one electron shell ends and the formation of the next, outer shell begins. In the Mendeleev system, this coincides with the beginning of a new period. Elements with 1, 2, 3, etc. electrons in a new shell are similar in their properties to those elements that also had 1, 2, 3, etc. outer electrons, although their number is inner. there were one (or several) fewer electron shells: Na is similar to Li (one external electron), Mg is similar to Be (2 external electrons); A1 - to B (3 external electrons), etc. With the position of the element in P. s. e. M. are associated with its chemical. and many more physical St.

Many (approx. 1000) graphic options have been proposed. images of P. s. e. M. The most common 2 variants of P. s. e. M. - short and long tables; k.-l. fundamental difference there is no between them. The appendix contains one of the short table options. In the table, the period numbers are given in the first column (indicated by Arabic numerals 1 - 7). The group numbers are indicated at the top with Roman numerals I - VIII. Each group is divided into two subgroups - a and b. A set of elements headed by elements of small periods, sometimes called. the main subgroups are a-m and (Li heads the subgroup alkali metals. F - halogens, He - inert gases etc.). In this case, the remaining subgroups of elements of large periods are called. side effects.

Elements with Z = 58 - 71 due to the special closeness of the structure of their atoms and the similarity of their chemistry. St. make up the lanthanide family, which is included in group III, but for convenience is placed at the bottom of the table. Elements with Z = 90 - 103 are often classified into the actinide family for the same reasons. They are followed by an element with Z = 104 - curchatovy and an element with Z = 105 (see Nilsborium). In July 1974 Owls. physicists reported the discovery of an element with Z = 106, and in Jan. 1976 - elements with Z = 107. Later elements with Z = 108 and 109 were synthesized. Lower. border of P. s. e. M. is known - it is given by hydrogen, since there cannot be an element with a nuclear charge less than one. The question is what upper limit P.S. e. M., i.e. to what extreme value art can reach. synthesis of elements remains unresolved. (Heavy nuclei are unstable, therefore americium with Z = 95 and subsequent elements are not found in nature, but are obtained in nuclear reactions; however, in the region of more distant transuranium elements, the appearance of the so-called. islands of stability, in particular for Z = 114.) In art. synthesis of new elements periodically. law and P. s. e. M. play a primary role. Mendeleev's law and system are among the most important generalizations of natural science and form the basis of modern science. teachings about the structure of the island.

Electronic structure of the atom.

This and the next paragraphs talk about models of the electron shell of an atom. It is important to understand that we're talking about exactly about models. Real atoms, of course, are more complex and we still don’t know everything about them. However, modern theoretical model the electronic structure of the atom makes it possible to successfully explain and even predict many properties of chemical elements, therefore it is widely used in the natural sciences.

To begin with, let us consider in more detail the “planetary” model proposed by N. Bohr (Fig. 2-3 c).

Rice. 2-3 c. Bohr's "planetary" model.

Danish physicist N. Bohr in 1913 proposed a model of the atom in which electron particles revolve around the atomic nucleus in approximately the same way as planets revolve around the Sun. Bohr suggested that electrons in an atom can exist stably only in orbits removed from the nucleus at strictly certain distances. He called these orbits stationary. Outside stationary orbits, an electron cannot exist. Why this was so, Bohr could not explain at that time. But he showed that such a model allows one to explain many experimental facts (this is discussed in more detail in paragraph 2.7).

Electron orbits in the Bohr model they are denoted by the integers 1, 2, 3, ... n, starting from the one closest to the core. In what follows we will call such orbits levels. To describe the electronic structure of the hydrogen atom, levels alone are sufficient. But more complex atoms, as it turned out, the levels consist of similar energies sublevels. For example, level 2 consists of two sublevels (2s and 2p). The third level consists of 3 sub-levels (3s, 3p and 3d), as shown in Fig. 2-6. The fourth level (it did not fit in the figure) consists of sublevels 4s, 4p, 4d, 4f. In paragraph 2.7 we will tell you where exactly these sublevel names came from and about physical experiments, which made it possible to “see” electronic levels and sublevels in atoms.

Rice. 2-6. Bohr's model for atoms more complex than the hydrogen atom. The drawing is not to scale - in fact, the sublevels of the same level are much closer closer friend to friend.

There are exactly as many electrons in the electron shell of any atom as there are protons in its nucleus, so the atom as a whole is electrically neutral. Electrons in an atom populate the levels and sublevels closest to the nucleus because in this case their energy is less than if they populated more distant levels. Each level and sublevel can only hold a certain number of electrons.

The sublevels, in turn, consist of equal energy orbitals(they are not shown in Fig. 2-6). Figuratively speaking, if the electron cloud of an atom is compared to a city or street where all the electrons of a given atom “live”, then a level can be compared to a house, a sublevel to an apartment, and an orbital to a room for electrons. All orbitals of any sublevel have the same energy. At the s-sublevel there is only one “room” - the orbital. The p-sublevel has 3 orbitals, the d-sublevel has 5, and the f-sublevel has as many as 7 orbitals. One or two electrons can “live” in each “room” orbital. The prohibition of electrons having more than two in one orbital is called Pauli's ban- named after the scientist who discovered this important feature structure of the atom. Each electron in an atom has its own "address", which is written as a set of four numbers called "quantum". Quantum numbers will be discussed in detail in section 2.7. Here we will mention only the main thing quantum number n(see Fig. 2-6), which in the “address” of the electron indicates the number of the level at which this electron exists.

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Everything in the world is made of atoms. But where did they come from, and what are they made of? Today we answer these simple and fundamental questions. After all, many people living on the planet say that they do not understand the structure of the atoms from which they themselves are composed.

Naturally, dear reader understands that in this article we try to present everything at the simplest and most interesting level, so we do not “load” scientific terms. For those who want to study the issue in more detail professional level, we recommend reading specialized literature. However, the information in this article may help good service in your studies and simply make you more erudite.

An atom is a particle of matter microscopic size and mass, the smallest part of a chemical element that is the carrier of its properties. In other words, this smallest particle a substance that can enter into chemical reactions.

Discovery history and structure

The concept of an atom was known back in Ancient Greece. Atomism – physical theory, which states that all material objects consist of indivisible particles. Along with Ancient Greece, the ideas of atomism also developed in parallel in Ancient India.

It is not known whether the aliens told the philosophers of that time about atoms, or whether they thought of it themselves, but it can be confirmed experimentally this theory chemists were able to do this much later - only in the seventeenth century, when Europe emerged from the abyss of the Inquisition and the Middle Ages.

For a long time, the dominant idea of ​​the structure of the atom was the idea of ​​it as an indivisible particle. The fact that the atom can still be divided became clear only at the beginning of the twentieth century. Rutherford, thanks to his famous experience with the deflection of alpha particles, learned that an atom consists of a nucleus around which electrons revolve. The planetary model of the atom was adopted, according to which electrons rotate around the nucleus, like our planets solar system around the star.

Modern representations much progress has been made about the structure of the atom. The nucleus of an atom, in turn, consists subatomic particles, or nucleons - protons and neutrons. It is nucleons that make up the bulk of the atom. Moreover, protons and neutrons are also not indivisible particles, and consist of fundamental particles - quarks.

The nucleus of an atom has a positive electric charge, and electrons rotating in orbit are negative. Thus, the atom is electrically neutral.

Below we give an elementary diagram of the structure of the carbon atom.

Properties of atoms

Weight

The mass of atoms is usually measured in atomic mass units - a.m.u. An atomic mass unit is the mass of 1/12 of a freely resting carbon atom in its ground state.

In chemistry, the concept is used to measure the mass of atoms "moth". 1 mole is the amount of substance that contains the number of atoms equal to the number Avogadro.

Size

The sizes of atoms are extremely small. So, the smallest atom is the Helium atom, its radius is 32 picometers. Most big atom– a cesium atom with a radius of 225 picometers. The prefix pico means ten to the minus twelfth power! That is, if we reduce 32 meters by a thousand billion times, we get the size of the radius of a helium atom.

At the same time, the scale of things is such that, in fact, the atom is 99% empty. The nucleus and electrons occupy an extremely small part of its volume. For clarity, consider this example. If you imagine an atom in the form of the Olympic stadium in Beijing (or maybe not in Beijing, just imagine a large stadium), then the nucleus of this atom will be a cherry located in the center of the field. The electron orbits would be somewhere at the level of the upper stands, and the cherry would weigh 30 million tons. Impressive, isn't it?

Where do atoms come from?

As you know, now various atoms grouped in the periodic table. There are 118 in it (and if with those predicted, but not yet open elements- 126) elements, not counting isotopes. But this was not always the case.

At the very beginning of the formation of the Universe, there were no atoms, and even more so, there were only elementary particles that interacted with each other under the influence of enormous temperatures. As a poet would say, it was a real apotheosis of particles. In the first three minutes of the existence of the Universe, due to a decrease in temperature and the coincidence of a whole bunch of factors, the process of primary nucleosynthesis began, when the first elements appeared from elementary particles: hydrogen, helium, lithium and deuterium (heavy hydrogen). It was from these elements that the first stars were formed, in the depths of which the thermonuclear reactions, as a result of which hydrogen and helium “burned”, forming heavier elements. If the star was large enough, then it ended its life with a so-called “supernova” explosion, as a result of which atoms were thrown into the surrounding space. This is how the entire periodic table turned out.

So, we can say that all the atoms that we are made of were once part of ancient stars.

Why doesn't the nucleus of an atom decay?

There are four types in physics fundamental interactions between particles and the bodies they compose. These are strong, weak, electromagnetic and gravitational interactions.

Thanks to strong interaction, which manifests itself on the scale of atomic nuclei and is responsible for the attraction between nucleons, the atom is such a “tough nut to crack.”

Not so long ago, people realized that when the nuclei of atoms split, enormous energy was released. The fission of heavy atomic nuclei is the source of energy in nuclear reactors and nuclear weapons.

So, friends, having introduced you to the structure and fundamentals of the structure of the atom, we can only remind you that we are ready to come to your aid at any time. It doesn’t matter, you need to complete a diploma in nuclear physics, or the smallest control - situations are different, but there is a way out of any situation. Think about the scale of the Universe, order work from Zaochnik and remember - there is no reason to worry.