DEFINITION

Oxygen- the eighth element of the Periodic Table. Refers to non-metals. Located in the second period of VI group A subgroup.

The serial number is 8. The nuclear charge is +8. Atomic weight - 15.999 amu. There are three isotopes of oxygen found in nature: 16 O, 17 O and 18 O, of which the most common is 16 O (99.762%).

Electronic structure of the oxygen atom

The oxygen atom has two shells, like all elements located in the second period. The group number -VI (chalcogens) - indicates that there are 6 valence electrons at the outer electronic level of the nitrogen atom. It has a high oxidizing ability (higher only for fluorine).

Rice. 1. Schematic representation of the structure of the oxygen atom.

The electronic configuration of the ground state is written as follows:

1s 2 2s 2 2p 4 .

Oxygen is an element of the p-family. The energy diagram for valence electrons in the unexcited state is as follows:

Oxygen has 2 pairs of paired electrons and two unpaired electrons. In all its compounds, oxygen exhibits valency II.

Rice. 2. Spatial representation of the structure of the oxygen atom.

Examples of problem solving

EXAMPLE 1

>> Chemistry: Electronic configurations of atoms of chemical elements

The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons having opposite (antiparallel) spins (translated from English as “spindle”), that is, having such properties that can be conventionally imagined itself as the rotation of an electron around its imaginary axis: clockwise or counterclockwise. This principle is called the Pauli principle.

If there is one electron in the orbital, then it is called unpaired; if there are two, then these are paired electrons, that is, electrons with opposite spins.

Figure 5 shows a diagram of the division of energy levels into sublevels.

The s-orbital, as you already know, has a spherical shape. The electron of the hydrogen atom (s = 1) is located in this orbital and is unpaired. Therefore, its electronic formula or electronic configuration will be written as follows: 1s 1. In electronic formulas, the number of the energy level is indicated by the number preceding the letter (1 ...), the Latin letter indicates the sublevel (type of orbital), and the number, which is written to the upper right of the letter (as an exponent), shows the number of electrons in the sublevel.

For a helium atom He, which has two paired electrons in one s-orbital, this formula is: 1s 2.

The electron shell of the helium atom is complete and very stable. Helium is a noble gas.

At the second energy level (n = 2) there are four orbitals: one s and three p. The electrons of the s-orbital of the second level (2s-orbitals) have higher energy, since they are at a greater distance from the nucleus than the electrons of the 1s-orbital (n = 2).

In general, for each value of n there is one s-orbital, but with a corresponding supply of electron energy on it and, therefore, with a corresponding diameter, growing as the value of n increases.

The p-Orbital has the shape of a dumbbell or a three-dimensional figure eight. All three p-orbitals are located in the atom mutually perpendicular along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized once again that each energy level (electronic layer), starting from n = 2, has three p-orbitals. As the value of n increases, electrons occupy p-orbitals located at large distances from the nucleus and directed along the x, y, z axes.

For elements of the second period (n = 2), first one b-orbital is filled, and then three p-orbitals. Electronic formula 1l: 1s 2 2s 1. The electron is more loosely bound to the nucleus of the atom, so the lithium atom can easily give it up (as you remember, this process is called oxidation), turning into a Li+ ion.

In the beryllium atom Be 0, the fourth electron is also located in the 2s orbital: 1s 2 2s 2. The two outer electrons of the beryllium atom are easily detached - Be 0 is oxidized into the Be 2+ cation.

In the boron atom, the fifth electron occupies the 2p orbital: 1s 2 2s 2 2p 1. Next, the C, N, O, E atoms are filled with 2p orbitals, which ends with the noble gas neon: 1s 2 2s 2 2p 6.

For elements of the third period, the Sv and Sr orbitals are filled, respectively. Five d-orbitals of the third level remain free:

11 Na 1s 2 2s 2 Sv1; 17С11в22822р63р5; 18Аг П^Ёр^Зр6.

Sometimes in diagrams depicting the distribution of electrons in atoms, only the number of electrons at each energy level is indicated, that is, abbreviated electronic formulas of atoms of chemical elements are written, in contrast to the full electronic formulas given above.

For elements of large periods (fourth and fifth), the first two electrons occupy the 4th and 5th orbitals, respectively: 19 K 2, 8, 8, 1; 38 Sr 2, 8, 18, 8, 2. Starting from the third element of each major period, the next ten electrons will enter the previous 3d and 4d orbitals, respectively (for elements of side subgroups): 23 V 2, 8, 11, 2; 26 Tr 2, 8, 14, 2; 40 Zr 2, 8, 18, 10, 2; 43 Tg 2, 8, 18, 13, 2. As a rule, when the previous d-sublevel is filled, the outer (4p- and 5p-respectively) p-sublevel will begin to fill.

For elements of large periods - the sixth and the incomplete seventh - electronic levels and sublevels are filled with electrons, as a rule, like this: the first two electrons will go to the outer b-sublevel: 56 Va 2, 8, 18, 18, 8, 2; 87Gg 2, 8, 18, 32, 18, 8, 1; the next one electron (for Na and Ac) to the previous one (p-sublevel: 57 La 2, 8, 18, 18, 9, 2 and 89 Ac 2, 8, 18, 32, 18, 9, 2.

Then the next 14 electrons will enter the third outer energy level in the 4f and 5f orbitals of the lanthanides and actinides, respectively.

Then the second external energy level (d-sublevel) will begin to build up again: for elements of secondary subgroups: 73 Ta 2, 8.18, 32.11, 2; 104 Rf 2, 8.18, 32, 32.10, 2, - and, finally, only after the current level is completely filled with ten electrons will the outer p-sublevel be filled again:

86 Rn 2, 8, 18, 32, 18, 8.

Very often, the structure of the electronic shells of atoms is depicted using energy or quantum cells - so-called graphical electronic formulas are written. For this notation, the following notation is used: each quantum cell is designated by a cell that corresponds to one orbital; Each electron is indicated by an arrow corresponding to the spin direction. When writing a graphical electronic formula, you should remember two rules: the Pauli principle, according to which there can be no more than two electrons in a cell (orbital), but with antiparallel spins, and F. Hund’s rule, according to which electrons occupy free cells (orbitals) and are located in At first, they are one at a time and have the same spin value, and only then they pair, but the spins will be oppositely directed according to the Pauli principle.

In conclusion, let us once again consider the display of the electronic configurations of atoms of elements according to the periods of the D.I. Mendeleev system. Diagrams of the electronic structure of atoms show the distribution of electrons across electronic layers (energy levels).

In a helium atom, the first electron layer is complete - it has 2 electrons.

Hydrogen and helium are s-elements; the s-orbital of these atoms is filled with electrons.

Elements of the second period

For all elements of the second period, the first electron layer is filled and electrons fill the e- and p-orbitals of the second electron layer in accordance with the principle of least energy (first s-, and then p) and the Pauli and Hund rules (Table 2).

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

Table 2 Structure of electronic shells of atoms of elements of the second period

End of table. 2

Li, Be - b-elements.

B, C, N, O, F, Ne are p-elements; these atoms have p-orbitals filled with electrons.

Elements of the third period

For atoms of elements of the third period, the first and second electronic layers are completed, so the third electronic layer is filled, in which electrons can occupy the 3s, 3p and 3d sublevels (Table 3).

Table 3 Structure of electronic shells of atoms of elements of the third period

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

An argon atom has 8 electrons in its outer layer (third electron layer). As an outer layer, it is complete, but in total in the third electron layer, as you already know, there can be 18 electrons, which means that the elements of the third period have unfilled 3d orbitals.

All elements from Al to Ar are p-elements. The s- and p-elements form the main subgroups in the Periodic Table.

A fourth electron layer appears in the potassium and calcium atoms, and the 4s sublevel is filled (Table 4), since it has lower energy than the 3d sublevel. To simplify the graphical electronic formulas of atoms of elements of the fourth period: 1) let us denote the conventional graphical electronic formula of argon as follows:
Ar;

2) we will not depict sublevels that are not filled in these atoms.

Table 4 Structure of electronic shells of atoms of elements of the fourth period

K, Ca - s-elements included in the main subgroups. In atoms from Sc to Zn, the 3rd sublevel is filled with electrons. These are Zy elements. They are included in secondary subgroups, their outermost electronic layer is filled, and they are classified as transition elements.

Pay attention to the structure of the electronic shells of chromium and copper atoms. In them there is a “failure” of one electron from the 4th to the 3rd sublevel, which is explained by the greater energy stability of the resulting electronic configurations Zd 5 and Zd 10:

In the zinc atom, the third electron layer is complete - all sublevels 3s, 3p and 3d are filled in it, with a total of 18 electrons.

In the elements following zinc, the fourth electron layer, the 4p sublevel, continues to be filled: Elements from Ga to Kr are p elements.

The krypton atom has an outer layer (fourth) that is complete and has 8 electrons. But in total in the fourth electron layer, as you know, there can be 32 electrons; the krypton atom still has unfilled 4d and 4f sublevels.

For elements of the fifth period, sublevels are filled in in the following order: 5s-> 4d -> 5p. And there are also exceptions associated with the “failure” of electrons in 41 Nb, 42 MO, etc.

In the sixth and seventh periods, elements appear, that is, elements in which the 4f- and 5f-sublevels of the third outside electronic layer are being filled, respectively.

4f elements are called lanthanides.

5f elements are called actinides.

The order of filling the electronic sublevels in the atoms of elements of the sixth period: 55 Сs and 56 Ва - 6s elements;

57 La... 6s 2 5d 1 - 5d element; 58 Ce - 71 Lu - 4f elements; 72 Hf - 80 Hg - 5d elements; 81 Tl- 86 Rn - 6p-elements. But here, too, there are elements in which the order of filling the electronic orbitals is “violated,” which, for example, is associated with the greater energy stability of half and completely filled f sublevels, that is, nf 7 and nf 14.

Depending on which sublevel of the atom is filled with electrons last, all elements, as you already understood, are divided into four electronic families or blocks (Fig. 7).

1) s-Elements; the b-sublevel of the outer level of the atom is filled with electrons; s-elements include hydrogen, helium and elements of the main subgroups of groups I and II;

2) p-elements; the p-sublevel of the outer level of the atom is filled with electrons; p elements include elements of the main subgroups of groups III-VIII;

3) d-elements; the d-sublevel of the pre-external level of the atom is filled with electrons; d-elements include elements of secondary subgroups of groups I-VIII, that is, elements of plug-in decades of large periods located between s- and p-elements. They are also called transition elements;

4) f-elements, the f-sublevel of the third outer level of the atom is filled with electrons; these include lanthanides and actinides.

1. What would happen if the Pauli principle were not observed?

2. What would happen if Hund's rule were not followed?

3. Make diagrams of the electronic structure, electronic formulas and graphic electronic formulas of atoms of the following chemical elements: Ca, Fe, Zr, Sn, Nb, Hf, Pa.

4. Write the electronic formula for element #110 using the appropriate noble gas symbol.

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The structure of the electronic shells of atoms of elements of the first four periods: $s-$, $p-$ and $d-$elements. Electronic configuration of an atom. Ground and excited states of atoms

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

Electrons

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, Mr. 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 a unit $(–1)$. Thomson even managed to determine the speed of the electron (it is equal to the speed of light - $300,000 km/s) and the mass of the electron (it is $1836$ times less than the mass of a hydrogen atom).

Thomson and Perrin connected the poles of a current source with two metal plates - a cathode and an anode, soldered into a glass tube from which the air was evacuated. When a voltage of about 10 thousand volts was applied to the electrode plates, a luminous discharge flashed in the tube, and particles flew from the cathode (negative pole) to the anode (positive pole), which scientists first called cathode rays, and then found out that it was a stream of electrons. Electrons hitting special substances, such as those on a TV screen, cause a glow.

The conclusion was drawn: electrons escape from the atoms of the material from which the cathode is made.

Free electrons or their flow can be obtained in other ways, for example, by heating a metal wire or by shining light on metals formed by elements of the main subgroup of group I of the periodic table (for example, cesium).

State of electrons in an atom

The state of an electron in an atom is understood as the totality of information about energy certain electron in space, in which it is located. We already know that an electron in an atom does not have a trajectory of motion, i.e. we can only talk about probabilities its location 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 a point. If countless such photographs were superimposed, the picture would be of an electron cloud with the greatest density where there are the most of these points.

The figure shows a “cut” of such an electron density in a hydrogen atom passing through the nucleus, and the dashed line outlines the sphere within which the probability of detecting an electron is $90%$. The contour closest to the nucleus covers a region of space in which the probability of detecting an electron is $10%$, the probability of detecting an electron inside the second contour from the nucleus is $20%$, inside the third - $≈30%$, etc. There is some uncertainty in the state of the electron. To characterize this special state, the German physicist W. Heisenberg introduced the concept of uncertainty principle, i.e. showed that it is impossible to simultaneously and accurately determine the energy and location of an electron. The more precisely the energy of an electron is determined, the more uncertain its position, and vice versa, having determined the position, it is impossible to determine the energy of the electron. The probability range for detecting an electron does not have clear boundaries. However, it is possible to select a space where the probability of finding an electron is maximum.

The space around the atomic nucleus in which an electron is most likely to be found is called an orbital.

It contains approximately $90%$ of the electron cloud, which means that about $90%$ of the time the electron is in this part of space. Based on their shape, there are four known types of orbitals, which are designated by the 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 the energy of its binding 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 $n$ denoting 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 are characterized by a large amount of energy. Consequently, the electrons of the outer level are least tightly bound to the atomic nucleus.

The number of energy levels (electronic layers) in an atom is equal to the number of the period in the D.I. Mendeleev system to which the chemical element belongs: atoms of elements of the first period have one energy level; second period - two; seventh period - seven.

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

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 $32$. And how, in turn, are energy levels (electronic layers) arranged?

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 sub level; the second - two; third - three; fourth - four. Sublevels, in turn, are formed by orbitals.

Each value of $n$ corresponds to a number of orbitals equal to $n^2$. According to the data presented in the table, one can trace the connection between the principal quantum number $n$ and the number of sublevels, the type and number of orbitals, and the maximum number of electrons at the sublevel and level.

Main quantum number, types and number of orbitals, maximum number of electrons in sublevels and levels.

Energy level $(n)$	Number of sublevels equal to $n$	Orbital type	Number of orbitals		Maximum number of electrons
			in the sublevel	in level equal to $n^2$	in the sublevel	at a level equal to $n^2$
$K(n=1)$	$1$	$1s$	$1$	$1$	$2$	$2$
$L(n=2)$	$2$	$2s$	$1$	$4$	$2$	$8$
		$2p$	$3$		$6$
$M(n=3)$	$3$	$3s$	$1$	$9$	$2$	$18$
		$3p$	$3$		$6$
		$3d$	$5$		$10$
$N(n=4)$	$4$	$4s$	$1$	$16$	$2$	$32$
		$4p$	$3$		$6$
		$4d$	$5$		$10$
		$4f$	$7$		$14$

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

• $s$-sublevel - the first sublevel of each energy level closest to the atomic nucleus, consists of one $s$-orbital;
• $p$-sublevel - the second sublevel of each, except the first, energy level, consists of three $p$-orbitals;
• $d$-sublevel - the third sublevel of each, starting from the third, energy level, consists of five $d$-orbitals;
• The $f$-sublevel of each, starting from the fourth energy level, consists of seven $f$-orbitals.

Atomic nucleus

But not only electrons are part of atoms. Physicist Henri Becquerel discovered that a natural mineral containing a uranium salt also emits unknown radiation, exposing photographic films shielded from light. This phenomenon was called radioactivity.

There are three types of radioactive rays:

1. $α$-rays, which consist of $α$-particles having a charge $2$ times greater than the charge of an electron, but with a positive sign, and a mass $4$ times greater than the mass of a hydrogen atom;
2. $β$-rays represent a flow of electrons;
3. $γ$-rays are electromagnetic waves with negligible mass that do not carry an electrical charge.

Consequently, the atom has a complex structure - it consists of a positively charged nucleus and electrons.

How is the atom structured?

In 1910, in Cambridge, near London, Ernest Rutherford and his students and colleagues studied the scattering of $α$ particles passing through thin gold foil and falling on a screen. The alpha particles usually deviated from the original direction by only one degree, seemingly confirming the uniformity and homogeneity of the properties of gold atoms. And suddenly the researchers noticed that some $α$ particles abruptly changed the direction of their path, as if encountering some kind of obstacle.

By placing a screen in front of the foil, Rutherford was able to detect even those rare cases when $α$ particles, reflected from gold atoms, flew in the opposite direction.

Calculations showed that the observed phenomena could occur if the entire mass of the atom and all its positive charge were concentrated in a tiny central nucleus. The radius of the nucleus, as it turned out, is 100,000 times smaller than the radius of the entire atom, the region in which electrons with a negative charge are located. If we apply a figurative comparison, then the entire volume of an atom can be likened to the stadium in Luzhniki, and the nucleus can be likened to a soccer ball located in the center of the field.

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

Protons and Neutrons

It turns out that the tiny 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 the electrons, but opposite in sign $(+1)$, and a mass equal to the mass of the hydrogen atom (it is taken as unity in chemistry). Protons are designated by the sign $↙(1)↖(1)p$ (or $p+$). Neutrons do not carry a charge, they are neutral and have a mass equal to the mass of a proton, i.e. $1$. Neutrons are designated by the sign $↙(0)↖(1)n$ (or $n^0$).

Protons and neutrons together are called nucleons(from lat. nucleus- core).

The sum of the number of protons and neutrons in an atom is called mass number. For example, the mass number of an aluminum atom is:

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 as follows: $e↖(-)$.

Since the atom is electrically neutral, it is also obvious that that the number of protons and electrons in an atom is the same. It is equal to the atomic number of the chemical element, assigned to it in the Periodic Table. For example, the nucleus of an iron atom contains $26$ protons, and $26$ electrons revolve around the nucleus. How to determine the number of neutrons?

As is known, the mass of an atom consists of the mass of protons and neutrons. Knowing the serial 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, the number of neutrons $(N)$ can be found using the formula:

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

$56 – 26 = 30$.

The table presents the main characteristics of elementary particles.

Basic characteristics of elementary particles.

Isotopes

Varieties of atoms of the same element that have the same nuclear charge but different mass numbers are called isotopes.

Word isotope consists of two Greek words: isos- identical and topos- place, means “occupying one place” (cell) in the Periodic Table of Elements.

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.

Usually, the relative atomic mass of a chemical element 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, therefore the values ​​of atomic masses are quite often fractional. For example, natural chlorine atoms are a mixture of two isotopes - $35$ (there are $75%$ in nature) and $37$ (they are $25%$ in nature); therefore, the relative atomic mass of chlorine is $35.5$. Isotopes of chlorine are written as follows:

$↖(35)↙(17)(Cl)$ and $↖(37)↙(17)(Cl)$

The chemical properties of chlorine isotopes are exactly the same, as are the isotopes of most chemical elements, for example potassium, argon:

$↖(39)↙(19)(K)$ and $↖(40)↙(19)(K)$, $↖(39)↙(18)(Ar)$ and $↖(40)↙(18 )(Ar)$

However, hydrogen isotopes vary greatly in properties due to the dramatic multiple increase in their relative atomic mass; they were even assigned individual names and chemical symbols: protium - $↖(1)↙(1)(H)$; deuterium - $↖(2)↙(1)(H)$, or $↖(2)↙(1)(D)$; tritium - $↖(3)↙(1)(H)$, or $↖(3)↙(1)(T)$.

Now we can give a modern, more rigorous and scientific definition of a chemical element.

A chemical element is a collection of atoms with the same nuclear charge.

The structure of the electronic shells of atoms of elements of the first four periods

Let's consider the display of electronic configurations of atoms of elements according to the periods of the D.I. Mendeleev system.

Elements of the first period.

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

Electronic formulas of atoms show the distribution of electrons across energy levels and sublevels.

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

In a helium atom, the first electron layer is complete - it contains $2$ electrons.

Hydrogen and helium are $s$ elements; the $s$ orbital of these atoms is filled with electrons.

Elements of the second period.

For all second-period elements, the first electron 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 contains $8$ electrons.

Elements of the third period.

For atoms of elements of the third period, the first and second electron layers are completed, so the third electron layer is filled, in which electrons can occupy the 3s-, 3p- and 3d-sub levels.

The structure of the electronic shells of atoms of elements of the third period.

The magnesium atom completes its $3.5$ electron orbital. $Na$ and $Mg$ are $s$-elements.

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

$↙(18)(Ar)$ Argon

$1s^2(2)s^2(2)p^6(3)s^2(3)p^6$

An argon atom has $8$ electrons in its outer layer (third electron layer). As the outer layer is completed, but in total in the third electron layer, as you already know, there can be 18 electrons, which means that the elements of the third period have $3d$ orbitals left unfilled.

All elements from $Al$ to $Ar$ are $р$ -elements.

$s-$ and $p$ -elements form main subgroups in the Periodic Table.

Elements of the fourth period.

Potassium and calcium atoms have a fourth electron layer and the $4s$ sublevel is filled, because it has lower energy than the $3d$ sublevel. To simplify the graphical electronic formulas of atoms of elements of the fourth period:

1. Let us denote the conventional graphical electronic formula of argon as follows: $Ar$;
2. We will not depict sublevels that are not filled in these atoms.

$K, Ca$ - $s$ -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 side subgroups, their outer electron layer is filled, they are classified as transitional 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 $3d^(10)$:

$↙(24)(Cr)$ $1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)3d^(4) 4s^(2)…$

$↙(29)(Cu)$ $1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)3d^(9)4s^(2)…$

Element symbol, serial number, name	Electronic structure diagram	Electronic formula	Graphical electronic formula
$↙(19)(K)$ Potassium		$1s^2(2)s^2(2)p^6(3)p^6(4)s^1$
$↙(20)(C)$ Calcium		$1s^2(2)s^2(2)p^6(3)p^6(4)s^2$
$↙(21)(Sc)$ Scandium		$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^1$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^1(4)s^1$
$↙(22)(Ti)$ Titanium		$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^2$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^2(4)s^2$
$↙(23)(V)$ Vanadium		$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^3$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^3(4)s^2$
$↙(24)(Cr)$ Chrome		$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^5$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^5(4)s^1$
$↙(29)(Cu)$ Chrome		$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^(10)$ or $1s^2(2)s^2(2 )p^6(3)p^6(3)d^(10)(4)s^1$
$↙(30)(Zn)$ Zinc		$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^(10)$ or $1s^2(2)s^2(2 )p^6(3)p^6(3)d^(10)(4)s^2$
$↙(31)(Ga)$ Gallium		$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^(10)4p^(1)$ or $1s^2(2) s^2(2)p^6(3)p^6(3)d^(10)(4)s^(2)4p^(1)$
$↙(36)(Kr)$ Krypton		$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^(10)4p^6$ or $1s^2(2)s^ 2(2)p^6(3)p^6(3)d^(10)(4)s^(2)4p^6$

In the zinc atom, the third electron layer is complete - all $3s, 3p$ and $3d$ sublevels 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 $Kr$ - $р$ -elements.

The outer (fourth) layer of the krypton atom is complete and has $8$ electrons. But in total in the fourth electron layer, as you know, there can be $32$ electrons; The krypton atom still has unfilled $4d-$ and $4f$ sublevels.

For elements of the fifth period, sublevels are filled in in the following order: $5s → 4d → 5p$. And there are also exceptions associated with the “failure” of electrons in $↙(41)Nb$, $↙(42)Mo$, $↙(44)Ru$, $↙(45)Rh$, $↙(46) Pd$, $↙(47)Ag$. $f$ appears in the sixth and seventh periods -elements, i.e. elements for which the $4f-$ and $5f$ sublevels of the third outside electronic layer are filled, respectively.

$4f$ -elements called lanthanides.

$5f$ -elements called actinides.

The order of filling electronic sublevels in atoms of elements of the sixth period: $↙(55)Cs$ and $↙(56)Ba$ - $6s$ elements; $↙(57)La ... 6s^(2)5d^(1)$ - $5d$-element; $↙(58)Се$ – $↙(71)Lu - 4f$-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 of electronic orbitals is violated, which, for example, is associated with greater energy stability of half and completely filled $f$-sublevels, i.e. $nf^7$ and $nf^(14)$.

Depending on which sublevel of the atom is filled with electrons last, all elements, as you already understood, are divided into four electron families, or blocks:

1. $s$ -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;
2. $p$ -elements; the $p$-sublevel of the outer level of the atom is filled with electrons; $p$-elements include elements of the main subgroups of groups III–VIII;
3. $d$ -elements; the $d$-sublevel of the pre-external level of the atom is filled with electrons; $d$-elements include elements of secondary subgroups of groups I–VIII, i.e. elements of intercalary decades of large periods located between $s-$ and $p-$elements. They are also called transition elements;
4. $f$ -elements; electrons fill the $f-$sublevel of the third outer level of the atom; these include lanthanides and actinides.

Electronic configuration of an atom. Ground and excited states of atoms

Swiss physicist W. Pauli in $1925 found that an atom can have no more than two electrons in one orbital, having opposite (antiparallel) backs (translated from English as a spindle), i.e. possessing properties that can be conventionally 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 an orbital, it is called unpaired, if two, then this paired electrons, i.e. electrons with opposite spins.

The figure shows a diagram of dividing energy levels into sublevels.

$s-$ Orbital, as you already know, has a spherical shape. The electron of the hydrogen atom $(n = 1)$ is located in this orbital and is unpaired. For this reason it electronic formula, or electronic configuration, is written like this: $1s^1$. In electronic formulas, the number of the energy level is indicated by the number in front of the letter $(1...)$, the Latin letter denotes the sublevel (type of orbital), and the number written to the right above the letter (as an exponent) shows the number of electrons in the sublevel.

For a helium atom He, which has two paired electrons in one $s-$orbital, this formula is: $1s^2$. The electron shell of the helium atom is complete and very stable. Helium is a noble gas. At the second energy level $(n = 2)$ there are four orbitals, one $s$ and three $p$. Electrons of the $s$-orbital of the second level ($2s$-orbital) have higher energy, because are at a greater distance from the nucleus than the electrons of the $1s$ orbital $(n = 2)$. In general, for each value of $n$ there is one $s-$orbital, but with a corresponding supply of electron energy on it and, therefore, with a corresponding diameter, growing as the value of $n$ increases. The $s-$Orbital, as you already know , has a spherical shape. The electron of the hydrogen atom $(n = 1)$ is located in this orbital and is unpaired. Therefore, its electronic formula, or electronic configuration, is written as follows: $1s^1$. In electronic formulas, the number of the energy level is indicated by the number in front of the letter $(1...)$, the Latin letter denotes the sublevel (type of orbital), and the number written to the right above the letter (as an exponent) shows the number of electrons in the sublevel.

For a helium atom $He$, which has two paired electrons in one $s-$orbital, this formula is: $1s^2$. The electron shell of the helium atom is complete and very stable. Helium is a noble gas. At the second energy level $(n = 2)$ there are four orbitals, one $s$ and three $p$. Electrons of $s-$orbitals of the second level ($2s$-orbitals) have higher energy, because are at a greater distance from the nucleus than the electrons of the $1s$ orbital $(n = 2)$. In general, for each value of $n$ there is one $s-$orbital, but with a corresponding supply of electron energy on it and, therefore, with a corresponding diameter, growing as the value of $n$ increases.

$p-$ Orbital has the shape of a dumbbell, or a voluminous figure eight. All three $p$-orbitals are located in the atom mutually perpendicular along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized once again that each energy level (electronic layer), starting from $n= 2$, has three $p$-orbitals. As the value of $n$ increases, electrons occupy $p$-orbitals located at large distances from the nucleus and directed along the $x, y, z$ axes.

For elements of the second period $(n = 2)$, first one $s$-orbital is filled, and then three $p$-orbitals; electronic formula $Li: 1s^(2)2s^(1)$. The $2s^1$ electron is more weakly bound to the nucleus of the atom, so the lithium atom can easily give it up (as you obviously remember, this process is called oxidation), turning into a lithium ion $Li^+$.

In the beryllium Be atom, the fourth electron is also located in the $2s$ orbital: $1s^(2)2s^(2)$. The two outer electrons of the beryllium atom are easily detached - $B^0$ is oxidized into the $Be^(2+)$ cation.

In the boron atom, the fifth electron occupies the $2p$ orbital: $1s^(2)2s^(2)2p^(1)$. Next, the $C, N, O, F$ atoms are filled with $2p$-orbitals, which ends with the noble gas neon: $1s^(2)2s^(2)2p^(6)$.

For elements of the third period, the $3s-$ and $3p$ orbitals are filled, respectively. Five $d$-orbitals of the third level remain free:

$↙(11)Na 1s^(2)2s^(2)2p^(6)3s^(1)$,

$↙(17)Cl 1s^(2)2s^(2)2p^(6)3s^(2)3p^(5)$,

$↙(18)Ar 1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)$.

Sometimes in diagrams depicting the distribution of electrons in atoms, only the number of electrons at each energy level is indicated, i.e. write abbreviated electronic formulas of atoms of chemical elements, in contrast to the full electronic formulas given above, for example:

$↙(11)Na 2, 8, 1;$ $↙(17)Cl 2, 8, 7;$ $↙(18)Ar 2, 8, 8$.

For elements of large periods (fourth and fifth), the first two electrons occupy $4s-$ and $5s$ orbitals, respectively: $↙(19)K 2, 8, 8, 1;$ $↙(38)Sr 2, 8, 18, 8, 2$. Starting from the third element of each major period, the next ten electrons will go to the previous $3d-$ and $4d-$orbitals, respectively (for elements of side subgroups): $↙(23)V 2, 8, 11, 2;$ $↙( 26)Fr 2, 8, 14, 2;$ $↙(40)Zr 2, 8, 18, 10, 2;$ $↙(43)Tc 2, 8, 18, 13, 2$. As a rule, when the previous $d$-sublevel is filled, the outer ($4р-$ and $5р-$, respectively) $р-$sublevel will begin to be filled: $↙(33)As 2, 8, 18, 5;$ $ ↙(52)Te 2, 8, 18, 18, 6$.

For elements of large periods - the sixth and incomplete seventh - electronic levels and sublevels are filled with electrons, as a rule, like this: the first two electrons enter the outer $s-$sublevel: $↙(56)Ba 2, 8, 18, 18, 8, 2;$ $↙(87)Fr 2, 8, 18, 32, 18, 8, 1$; the next one electron (for $La$ and $Ca$) to the previous $d$-sublevel: $↙(57)La 2, 8, 18, 18, 9, 2$ and $↙(89)Ac 2, 8, 18, 32, 18, 9, 2$.

Then the next $14$ electrons will go to the third outer energy level, to the $4f$ and $5f$ orbitals of lanthanides and actinides, respectively: $↙(64)Gd 2, 8, 18, 25, 9, 2;$ $↙(92 )U 2, 8, 18, 32, 21, 9, 2$.

Then the second external energy level ($d$-sublevel) of elements of side subgroups will begin to build up again: $↙(73)Ta 2, 8, 18, 32, 11, 2;$ $↙(104)Rf 2, 8, 18 , 32, 32, 10, 2$. And finally, only after the $d$-sublevel is completely filled with ten electrons will the $p$-sublevel be filled again: $↙(86)Rn 2, 8, 18, 32, 18, 8$.

Very often the structure of the electronic shells of atoms is depicted using energy or quantum cells - the so-called graphic electronic formulas. 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 principle, according to which there can be no more than two electrons in a cell (orbital), but with antiparallel spins, and F. Hund's rule, according to which electrons occupy free cells first one at a time and have the same spin value, and only then pair, but the spins, according to the Pauli principle, will be in opposite directions.

The filling of orbitals in a non-excited atom is carried out in such a way that the energy of the atom is minimal (the principle of minimum energy). First, the orbitals of the first energy level are filled, then the second, and the orbital of the s-sublevel is filled first and only then the orbitals of the p-sublevel. In 1925, the Swiss physicist W. Pauli established the fundamental quantum mechanical principle of natural science (the Pauli principle, also called the exclusion principle or the exclusion principle). According to the Pauli principle:

An atom cannot have two electrons that have the same set of all four quantum numbers.

The electronic configuration of an atom is expressed by a formula in which the filled orbitals are indicated by a combination of a number equal to the principal quantum number and a letter corresponding to the orbital quantum number. The superscript indicates the number of electrons in these orbitals.

Hydrogen and helium

The electronic configuration of the hydrogen atom is 1s 1, and the helium atom is 1s 2. A hydrogen atom has one unpaired electron, and a helium atom has two paired electrons. Paired electrons have the same values ​​of all quantum numbers except the spin one. A hydrogen atom can give up its electron and turn into a positively charged ion - the H + cation (proton), which has no electrons (electronic configuration 1s 0). A hydrogen atom can add one electron and become a negatively charged H - ion (hydride ion) with the electron configuration 1s 2.

Lithium

The three electrons in a lithium atom are distributed as follows: 1s 2 1s 1. Only electrons from the outer energy level, called valence electrons, participate in the formation of a chemical bond. In a lithium atom, the valence electron is the 2s sublevel electron, and the two electrons of the 1s sublevel are internal electrons. The lithium atom quite easily loses its valence electron, transforming into the Li + ion, which has the 1s 2 2s 0 configuration. Note that the hydride ion, helium atom, and lithium cation have the same number of electrons. Such particles are called isoelectronic. They have similar electronic configurations but different nuclear charges. The helium atom is very chemically inert, which is due to the special stability of the 1s 2 electronic configuration. Orbitals that are not filled with electrons are called vacant. In the lithium atom, three orbitals of the 2p sublevel are vacant.

Beryllium

The electronic configuration of the beryllium atom is 1s 2 2s 2. When an atom is excited, electrons from a lower energy sublevel move to vacant orbitals of a higher energy sublevel. The process of excitation of a beryllium atom can be conveyed by the following diagram:

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

A comparison of the ground and excited states of the beryllium atom shows that they differ in the number of unpaired electrons. In the ground state of the beryllium atom there are no unpaired electrons; in the excited state there are two. Despite the fact that when an atom is excited, in principle, any electrons from lower energy orbitals can move to higher orbitals, for the consideration of chemical processes only transitions between energy sublevels with similar energies are significant.

This is explained as follows. When a chemical bond is formed, energy is always released, i.e., the combination of two atoms transforms into an energetically more favorable state. The process of excitation requires energy expenditure. When pairing electrons within the same energy level, the excitation costs are compensated by the formation of a chemical bond. When pairing electrons within different levels, the excitation costs are so high that they cannot be compensated by the formation of a chemical bond. In the absence of a partner in a possible chemical reaction, the excited atom releases a quantum of energy and returns to the ground state - this process is called relaxation.

Bor

The electronic configurations of atoms of elements of the 3rd period of the Periodic Table of Elements will be to a certain extent similar to those given above (the subscript indicates the atomic number):

11 Na 3s 1
12 Mg 3s 2
13 Al 3s 2 3p 1
14 Si 2s 2 2p2
15P 2s 2 3p 3

However, the analogy is not complete, since the third energy level is split into three sublevels and all of the listed elements have vacant d-orbitals to which electrons can transfer upon excitation, increasing multiplicity. This is especially important for elements such as phosphorus, sulfur and chlorine.

The maximum number of unpaired electrons in a phosphorus atom can reach five:

This explains the possibility of the existence of compounds in which the valence of phosphorus is 5. The nitrogen atom, which has the same configuration of valence electrons in the ground state as the phosphorus atom, cannot form five covalent bonds.

A similar situation arises when comparing the valence capabilities of oxygen and sulfur, fluorine and chlorine. The pairing of electrons in a sulfur atom results in the appearance of six unpaired electrons:

3s 2 3p 4 (ground state) → 3s 1 3p 3 3d 2 (excited state).

This corresponds to the six-valence state, which is unattainable for oxygen. The maximum valency of nitrogen (4) and oxygen (3) requires a more detailed explanation, which will be given later.

The maximum valency of chlorine is 7, which corresponds to the configuration of the excited state of the atom 3s 1 3p 3 d 3.

The presence of vacant 3d orbitals in all elements of the third period is explained by the fact that, starting from the 3rd energy level, partial overlap of sublevels of different levels occurs when filled with electrons. Thus, the 3d sublevel begins to fill only after the 4s sublevel is filled. The energy reserve of electrons in atomic orbitals of different sublevels and, consequently, the order of their filling increases in the following order:

Orbitals for which the sum of the first two quantum numbers (n + l) is smaller are filled earlier; if these sums are equal, the orbitals with the lower principal quantum number are filled first.

This pattern was formulated by V. M. Klechkovsky in 1951.

Elements in whose atoms the s-sublevel is filled with electrons are called s-elements. These include the first two elements of each period: hydrogen. However, already in the next d-element - chromium - there is some “deviation” in the arrangement of electrons in energy levels in the ground state: instead of the expected four unpaired electrons on the 3d sublevel, the chromium atom has five unpaired electrons in the 3d sublevel and one unpaired electron in the s sublevel: 24 Cr 4s 1 3d 5 .

The phenomenon of the transition of one s-electron to the d-sublevel is often called “leakthrough” of an electron. This can be explained by the fact that the orbitals of the d-sublevel filled by electrons become closer to the nucleus due to increased electrostatic attraction between electrons and the nucleus. As a result, the state 4s 1 3d 5 becomes energetically more favorable than 4s 2 3d 4. Thus, the half-filled d-sublevel (d 5) has increased stability compared to other possible electron distribution options. The electronic configuration corresponding to the existence of the maximum possible number of paired electrons, achievable in previous d-elements only as a result of excitation, is characteristic of the ground state of the chromium atom. The electronic configuration d 5 is also characteristic of the manganese atom: 4s 2 3d 5. For the following d-elements, each energy cell of the d-sublevel is filled with a second electron: 26 Fe 4s 2 3d 6 ; 27 Co 4s 2 3d 7 ; 28 Ni 4s 2 3d 8 .

In the copper atom, the state of a completely filled d-sublevel (d 10) becomes achievable due to the transition of one electron from the 4s sub-level to the 3d sublevel: 29 Cu 4s 1 3d 10. The last element of the first row of d-elements has the electronic configuration 30 Zn 4s 23 d 10.

The general trend, manifested in the stability of the d 5 and d 10 configurations, is also observed in elements of lower periods. Molybdenum has an electronic configuration similar to chromium: 42 Mo 5s 1 4d 5, and silver to copper: 47 Ag5s 0 d 10. Moreover, the d 10 configuration is already achieved in palladium due to the transition of both electrons from the 5s orbital to the 4d orbital: 46Pd 5s 0 d 10. There are other deviations from the monotonic filling of d- and f-orbitals.

DEFINITION

Fluorine- an element belonging to the halogen group. Non-metal. Located in the second period of VII group A subgroup.

The serial number is 9. The nuclear charge is +9. Atomic weight - 18.998 amu. It is the only stable fluorine nuclide.

Electronic structure of the fluorine atom

The fluorine atom has two shells, like all elements located in the second period. The group number - VII (halogens) - indicates that the outer electronic level of the nitrogen atom contains 7 valence electrons and only one electron is missing to complete the outer energy level. It has the highest oxidizing capacity among all elements of the Periodic Table.

Rice. 1. Conventional representation of the structure of the fluorine atom.

The electronic configuration of the ground state is written as follows:

1s 2 2s 2 2p 5 .

Fluorine is an element of the p-family. The energy diagram for valence electrons in the unexcited state is as follows:

Fluorine has 3 pairs of paired electrons and one unpaired electron. In all its compounds, fluorine exhibits valency I and oxidation state -1.

As a result of the interaction, fluorine becomes an electron acceptor. In this case, the atom turns into a negatively charged ion (F -).