Electronic structure of the atom table. Structure and principles of the atom

Composition of the atom.

An atom is made up of atomic nucleus And electron shell.

The nucleus of an atom consists of protons ( p+) and neutrons ( n 0). Most hydrogen atoms have a nucleus consisting of one proton.

Number of protons N(p+) is equal to the nuclear charge ( Z) and the ordinal number of the element in the natural series of elements (and in the periodic table of elements).

N(p +) = Z

Sum of neutrons N(n 0), denoted simply by the letter N, and number of protons Z called mass number and is designated by the letter A.

A = Z + N

The electron shell of an atom consists of electrons moving around the nucleus ( e -).

Number of electrons N(e-) in the electron shell of a neutral atom is equal to the number of protons Z at its core.

The mass of a proton is approximately equal to the mass of a neutron and 1840 times the mass of an electron, so the mass of an atom is almost equal to the mass of the nucleus.

The shape of the atom is spherical. The radius of the nucleus is approximately 100,000 times smaller than the radius of the atom.

Chemical element- type of atoms (collection of atoms) with the same nuclear charge (with the same number of protons in the nucleus).

Isotope- a collection of atoms of the same element with the same number of neutrons in the nucleus (or a type of atom with the same number of protons and the same number of neutrons in the nucleus).

Different isotopes differ from each other in the number of neutrons in the nuclei of their atoms.

Designation of an individual atom or isotope: (E - element symbol), for example: .

Structure of the electron shell of an atom

Atomic orbital- state of an electron in an atom. The symbol for the orbital is . Each orbital has a corresponding electron cloud.

Orbitals of real atoms in the ground (unexcited) state are of four types: s, p, d And f.

Electronic cloud- the part of space in which an electron can be found with a probability of 90 (or more) percent.

Note: sometimes the concepts of “atomic orbital” and “electron cloud” are not distinguished, calling both “atomic orbital”.

The electron shell of an atom is layered. Electronic layer formed by electron clouds of the same size. The orbitals of one layer form electronic ("energy") level, their energies are the same for the hydrogen atom, but different for other atoms.

Orbitals of the same type are grouped into electronic (energy) sublevels:
s-sublevel (consists of one s-orbitals), symbol - .
p-sublevel (consists of three p
d-sublevel (consists of five d-orbitals), symbol - .
f-sublevel (consists of seven f-orbitals), symbol - .

The energies of orbitals of the same sublevel are the same.

When designating sublevels, the number of the layer (electronic level) is added to the sublevel symbol, for example: 2 s, 3p, 5d means s-sublevel of the second level, p-sublevel of the third level, d-sublevel of the fifth level.

The total number of sublevels at one level is equal to the level number n. The total number of orbitals at one level is n 2. Accordingly, the total number of clouds in one layer is also equal to n 2 .

Designations: - free orbital (without electrons), - orbital with an unpaired electron, - orbital with an electron pair (with two electrons).

The order in which electrons fill the orbitals of an atom is determined by three laws of nature (the formulations are given in simplified terms):

1. The principle of least energy - electrons fill the orbitals in order of increasing energy of the orbitals.

2. The Pauli principle - there cannot be more than two electrons in one orbital.

3. Hund's rule - within a sublevel, electrons first fill empty orbitals (one at a time), and only after that they form electron pairs.

The total number of electrons in the electronic level (or electron layer) is 2 n 2 .

The distribution of sublevels by energy is expressed as follows (in order of increasing energy):

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

This sequence is clearly expressed by an energy diagram:

The distribution of an atom's electrons across levels, sublevels, and orbitals (electronic configuration of an atom) can be depicted as an electron formula, an energy diagram, or, more simply, as a diagram of electron layers ("electron diagram").

Examples of the electronic structure of atoms:

Valence electrons- electrons of an atom that can take part in the formation of chemical bonds. For any atom, these are all the outer electrons plus those pre-outer electrons whose energy is greater than that of the outer ones. For example: the Ca atom has 4 outer electrons s 2, they are also valence; the Fe atom has 4 outer electrons s 2 but he has 3 d 6, therefore the iron atom has 8 valence electrons. Valence electronic formula of the calcium atom is 4 s 2, and iron atoms - 4 s 2 3d 6 .

Periodic table of chemical elements by D. I. Mendeleev
(natural system of chemical elements)

Periodic law of chemical elements(modern formulation): the properties of chemical elements, as well as simple and complex substances formed by them, are periodically dependent on the value of the charge of atomic nuclei.

Periodic table- graphic expression of the periodic law.

Natural series of chemical elements- a series of chemical elements arranged according to the increasing number of protons in the nuclei of their atoms, or, what is the same, according to the increasing charges of the nuclei of these atoms. The atomic number of an element in this series is equal to the number of protons in the nucleus of any atom of this element.

The table of chemical elements is constructed by "cutting" the natural series of chemical elements into periods(horizontal rows of the table) and groupings (vertical columns of the table) of elements with a similar electronic structure of atoms.

Depending on the way elements are grouped, the table may be long-period(elements with the same number and type of valence electrons are collected into groups) and short period(elements with the same number of valence electrons are collected into groups).

The short-period table groups are divided into subgroups ( main And side), coinciding with the groups of the long-period table.

All atoms of elements of the same period have the same number of electron layers, equal to the period number.

Number of elements in periods: 2, 8, 8, 18, 18, 32, 32. Most of the elements of the eighth period were obtained artificially; the last elements of this period have not yet been synthesized. All periods except the first begin with an alkali metal-forming element (Li, Na, K, etc.) and end with a noble gas-forming element (He, Ne, Ar, Kr, etc.).

In the short-period table there are eight groups, each of which is divided into two subgroups (main and secondary), in the long-period table there are sixteen groups, which are numbered in Roman numerals with the letters A or B, for example: IA, IIIB, VIA, VIIB. Group IA of the long-period table corresponds to the main subgroup of the first group of the short-period table; group VIIB - secondary subgroup of the seventh group: the rest - similarly.

The characteristics of chemical elements naturally change in groups and periods.

In periods (with increasing serial number)

• nuclear charge increases
• the number of outer electrons increases,
• the radius of atoms decreases,
• the strength of the bond between electrons and the nucleus increases (ionization energy),
• electronegativity increases,
• the oxidizing properties of simple substances are enhanced ("non-metallicity"),
• the reducing properties of simple substances weaken ("metallicity"),
• weakens the basic character of hydroxides and corresponding oxides,
• the acidic character of hydroxides and corresponding oxides increases.

In groups (with increasing serial number)

• nuclear charge increases
• the radius of atoms increases (only in A-groups),
• the strength of the bond between electrons and the nucleus decreases (ionization energy; only in A-groups),
• electronegativity decreases (only in A-groups),
• the oxidizing properties of simple substances weaken ("non-metallicity"; only in A-groups),
• the reducing properties of simple substances are enhanced ("metallicity"; only in A-groups),
• the basic character of hydroxides and corresponding oxides increases (only in A-groups),
• weakens the acidic character of hydroxides and corresponding oxides (only in A-groups),
• the stability of hydrogen compounds decreases (their reducing activity increases; only in A-groups).

Tasks and tests on the topic "Topic 9. "Structure of the atom. Periodic law and periodic system of chemical elements by D. I. Mendeleev (PSHE) "."

• Periodic law - Periodic law and structure of atoms grades 8–9
  You must know: the laws of filling orbitals with electrons (the principle of least energy, the Pauli principle, Hund's rule), the structure of the periodic table of elements.

  You must be able to: determine the composition of an atom by the position of the element in the periodic table, and, conversely, find an element in the periodic system, knowing its composition; depict the structure diagram, electronic configuration of an atom, ion, and, conversely, determine the position of a chemical element in the PSCE from the diagram and electronic configuration; characterize the element and the substances it forms according to its position in the PSCE; determine changes in the radius of atoms, properties of chemical elements and the substances they form within one period and one main subgroup of the periodic system.

  Example 1. Determine the number of orbitals in the third electron level. What are these orbitals?
  To determine the number of orbitals, we use the formula N orbitals = n 2 where n- level number. N orbitals = 3 2 = 9. One 3 s-, three 3 p- and five 3 d-orbitals.

  Example 2. Determine which element's atom has electronic formula 1 s 2 2s 2 2p 6 3s 2 3p 1 .
  In order to determine what element it is, you need to find out its atomic number, which is equal to the total number of electrons of the atom. In this case: 2 + 2 + 6 + 2 + 1 = 13. This is aluminum.

  After making sure that everything you need has been learned, proceed to completing the tasks. We wish you success.

  
  Recommended reading:
  • O. S. Gabrielyan and others. Chemistry 11th grade. M., Bustard, 2002;
  • G. E. Rudzitis, F. G. Feldman. Chemistry 11th grade. M., Education, 2001.

Atom- the smallest particle of a substance that is indivisible by chemical means. In the 20th century, the complex structure of the atom was discovered. Atoms are made up of positively charged kernels and a shell formed by negatively charged electrons. The total charge of a free atom is zero, since the charges of the nucleus and electron shell balance each other. In this case, the nuclear charge is equal to the number of the element in the periodic table ( atomic number) and is equal to the total number of electrons (electron charge is −1).

The atomic nucleus consists of positively charged protons and neutral particles - neutrons, having no charge. Generalized characteristics of elementary particles in an atom can be presented in the form of a table:

The number of protons is equal to the charge of the nucleus, therefore equal to the atomic number. To find the number of neutrons in an atom, you need to subtract the charge of the nucleus (the number of protons) from the atomic mass (consisting of the masses of protons and neutrons).

For example, in the sodium atom 23 Na the number of protons is p = 11, and the number of neutrons is n = 23 − 11 = 12

The number of neutrons in atoms of the same element can be different. Such atoms are called isotopes .

The electron shell of an atom also has a complex structure. Electrons are located in energy levels (electronic layers).

The level number characterizes the energy of the electron. This is due to the fact that elementary particles can transmit and receive energy not in arbitrarily small quantities, but in certain portions - quanta. The higher the level, the more energy the electron has. Since the lower the energy of the system, the more stable it is (compare the low stability of a stone on top of a mountain, which has high potential energy, and the stable position of the same stone below on the plain, when its energy is much lower), the levels with low electron energy are filled first and only then - high.

The maximum number of electrons that a level can accommodate can be calculated using the formula:
N = 2n 2, where N is the maximum number of electrons at the level,
n - level number.

Then for the first level N = 2 1 2 = 2,

for the second N = 2 2 2 = 8, etc.

The number of electrons in the outer level for elements of the main (A) subgroups is equal to the group number.

In most modern periodic tables, the arrangement of electrons by level is indicated in the cell with the element. Very important understand that the levels are readable from bottom to top, which corresponds to their energy. Therefore, the column of numbers in the cell with sodium:
1
8
2

at the 1st level - 2 electrons,

at the 2nd level - 8 electrons,

at the 3rd level - 1 electron
Be careful, this is a very common mistake!

The electron level distribution can be represented as a diagram:
11 Na)))
2 8 1

If the periodic table does not indicate the distribution of electrons by level, you can use:

• maximum number of electrons: at the 1st level no more than 2 e − ,
  on the 2nd - 8 e − ,
  at the external level - 8 e − ;
• number of electrons in the outer level (for the first 20 elements coincides with the group number)

Then for sodium the line of reasoning will be as follows:

1. The total number of electrons is 11, therefore, the first level is filled and contains 2 e − ;
2. The third, outer level contains 1 e − (I group)
3. The second level contains the remaining electrons: 11 − (2 + 1) = 8 (completely filled)

* A number of authors, in order to more clearly distinguish between a free atom and an atom in a compound, propose to use the term “atom” only to designate a free (neutral) atom, and to designate all atoms, including those in compounds, propose the term “atomic particles”. Time will tell what the fate of these terms will be. From our point of view, an atom by definition is a particle, therefore, the expression “atomic particles” can be considered as a tautology (“oil”).

2. Task. Calculation of the amount of substance of one of the reaction products if the mass of the starting substance is known.
Example:

What amount of hydrogen substance will be released when zinc reacts with hydrochloric acid weighing 146 g?

Solution:

1. We write the reaction equation: Zn + 2HCl = ZnCl 2 + H 2
2. Find the molar mass of hydrochloric acid: M (HCl) = 1 + 35.5 = 36.5 (g/mol)
   (the molar mass of each element, numerically equal to the relative atomic mass, is looked at in the periodic table under the sign of the element and rounded to whole numbers, except for chlorine, which is taken as 35.5)
3. Find the amount of hydrochloric acid: n (HCl) = m / M = 146 g / 36.5 g/mol = 4 mol
4. We write down the available data above the reaction equation, and below the equation - the number of moles according to the equation (equal to the coefficient in front of the substance):
   4 mol x mol
   Zn + 2HCl = ZnCl 2 + H 2
   2 mole 1 mole
5. Let's make a proportion:
   4 mol - x mole
   2 mol - 1 mol
   (or with an explanation:
   from 4 moles of hydrochloric acid you get x mole of hydrogen,
   and from 2 moles - 1 mole)
6. We find x:
   x= 4 mol 1 mol / 2 mol = 2 mol

Answer: 2 mol.

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, the modulus is equal to the charge of a proton, but 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 Russian chemist D.I. Mendeleev, having discovered the periodic law and developed the 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 the chemical properties of the atom, but more on that 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 mass unit (amu) is a special physical quantity to designate 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 of 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 is 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 we take the classical definition from quantum mechanics, then spin is the particle’s own angular momentum. 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 indicated by an arrow: up arrow – electron with positive spin +½, down arrow ↓ – electron with 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 of the same shape form a sublevel. The number of orbitals at a sublevel is determined by the number of possible locations 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.

The d-orbital has a complex shape:

In space, the d-orbital can be positioned 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 the school course, f-orbitals are not filled, so we don’t have to depict 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. An electronic formula is a short one-line representation of the distribution of electrons among 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. In case of failure, one electron from the s-sublevel of the next energy level literally falls into a d-electron.

Ground and excited states of the atom.

The atoms whose electronic configurations we have now constructed 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 elements of group VIII-A (the eighth group of the main subgroup), as stated in the next paragraph, are noble gases, so 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.

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 orbital number n. The radii of the orbits increase in proportion to the square of the orbital 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 electron 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 electron 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 in opposite directions.

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 first filled, 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

The concept of “atom” has been familiar to humanity since the times of Ancient Greece. According to the statement of ancient philosophers, an atom is the smallest particle that is part of a substance.

Electronic structure of the atom

An atom consists of a positively charged nucleus containing protons and neutrons. Electrons move in orbits around the nucleus, each of which can be characterized by a set of four quantum numbers: principal (n), orbital (l), magnetic (ml) and spin (ms or s).

The principal quantum number determines the energy of the electron and the size of the electron clouds. The energy of an electron mainly depends on the distance of the electron from the nucleus: the closer the electron is to the nucleus, the lower its energy. In other words, the principal quantum number determines the location of the electron at a particular energy level (quantum layer). The principal quantum number has the values ​​of a series of integers from 1 to infinity.

The orbital quantum number characterizes the shape of the electron cloud. The different shapes of electron clouds cause a change in the energy of electrons within one energy level, i.e. splitting it into energy sublevels. The orbital quantum number can have values ​​from zero to (n-1), for a total of n values. Energy sublevels are designated by letters:

The magnetic quantum number shows the orientation of the orbital in space. It accepts any integer value from (+l) to (-l), including zero. The number of possible values ​​of the magnetic quantum number is (2l+1).

An electron, moving in the field of the atomic nucleus, in addition to the orbital angular momentum, also has its own angular momentum, which characterizes its spindle-shaped rotation around its own axis. This property of an electron is called spin. The magnitude and orientation of the spin is characterized by the spin quantum number, which can take values ​​(+1/2) and (-1/2). Positive and negative spin values ​​are related to its direction.

Before all of the above became known and confirmed experimentally, there were several models of the structure of the atom. One of the first models of the structure of the atom was proposed by E. Rutherford, who, in experiments on the scattering of alpha particles, showed that almost the entire mass of the atom is concentrated in a very small volume - a positively charged nucleus. According to his model, electrons move around the nucleus at a sufficiently large distance, and their number is such that, on the whole, the atom is electrically neutral.

Rutherford's model of the structure of the atom was developed by N. Bohr, who in his research also combined Einstein's teachings on light quanta and Planck's quantum theory of radiation. Louis de Broglie and Schrödinger completed what they started and presented to the world a modern model of the structure of the atom of a chemical element.

Examples of problem solving

EXAMPLE 1

EXAMPLE 2