How to determine the electronic structure of an atom. The structure of atoms of chemical elements

All matter is made up of very small particles called atoms . An atom is the smallest particle of a chemical element that retains all of its characteristic properties. To imagine the size of an atom, it is enough to say that if they could be placed close to each other, then one million atoms would occupy a distance of only 0.1 mm.

Further development of the science of the structure of matter showed that the atom also has a complex structure and consists of electrons and protons. This is how the electronic theory of the structure of matter arose.

In ancient times it was discovered that there are two kinds of electricity: positive and negative. The amount of electricity contained in the body came to be called charge. Depending on the kind of electricity that the body possesses, the charge can be positive or negative.

It was also established empirically that like charges repel each other, and opposite charges attract.

Consider electronic structure of the atom. Atoms are made up of even smaller particles than themselves, called electrons.

DEFINITION:An electron is the smallest particle of matter that has the smallest negative electric charge.

Electrons revolve around a central nucleus made up of one or more protons and neutrons, in concentric orbits. Electrons are negatively charged particles, protons are positive, and neutrons are neutral (Figure 1.1).

DEFINITION:Proton is the smallest particle of matter that has the smallest positive electric charge.

The existence of electrons and protons is beyond doubt. Scientists not only determined the mass, charge and size of electrons and protons, but even made them work in various electrical and radio engineering devices.

It was also found that the mass of an electron depends on the speed of its movement and that the electron not only moves forward in space, but also rotates around its axis.

The simplest in its structure is the hydrogen atom (Fig. 1.1). It consists of a proton nucleus and an electron rotating around the nucleus at great speed, forming the outer shell (orbit) of the atom. More complex atoms have multiple shells around which electrons revolve.

These shells are sequentially filled with electrons from the nucleus (Figure 1.2).

Now let's analyze . The outermost shell is called valence, and the number of electrons it contains is called valence. The further away from the core valence shell, consequently, the smaller the force of attraction experienced by each valence electron from the side of the nucleus. Thus, the atom increases the ability to attach electrons to itself if the valence shell is not filled and is located far from the nucleus, or lose them.
The outer shell electrons can receive energy. If the electrons in the valence shell receive the required level of energy from external forces, they can break away from it and leave the atom, that is, become free electrons. Free electrons are able to move arbitrarily from one atom to atom. Those materials that contain a large number of free electrons are called conductors .

insulators , is the opposite of conductors. They block the flow of electricity. Insulators are stable because the valence electrons of some atoms fill the valence shells of other atoms, joining them. This prevents the formation of free electrons.
An intermediate position between insulators and conductors is occupied by semiconductors but we'll talk about them later.
Consider properties of an atom. An atom that has the same number of electrons and protons is electrically neutral. An atom that receives one or more electrons becomes negatively charged and is called a negative ion. If an atom loses one or more electrons, then it becomes a positive ion, that is, it becomes positively charged.

The concept of an atom arose in the ancient world to designate the particles of matter. In Greek, atom means "indivisible".

Electrons

The Irish physicist Stoney, on the basis of experiments, came to the conclusion that electricity is carried by the smallest particles that exist in the atoms of all chemical elements. In $1891$, Stoney proposed to call these particles electrons, which in Greek means "amber".

A few years after the electron got its name, English physicist Joseph Thomson and French physicist Jean Perrin proved that electrons carry a negative charge. This is the smallest negative charge, which in chemistry is taken as the 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 the 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 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 applied, for example, to a TV screen, cause a glow.

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

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

The state of electrons in an atom

The state of an electron in an atom is understood as a set of information about energy specific 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. can only talk about probabilities 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 as follows: if it were possible to photograph the position of an electron in an atom in hundredths or millionths of a second, as in a photo finish, then the electron in such photographs would be represented as a point. Overlaying countless such photographs would result in a picture of an electron cloud with the highest density where there are 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 delimits the sphere within which the probability of finding an electron is $90%$. The contour closest to the nucleus covers the region of space in which the probability of finding an electron is $10%$, the probability of finding an electron inside the second contour from the nucleus is $20%$, inside the third one - $≈30%$, 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 determine simultaneously and exactly the energy and location of the electron. The more accurately 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 electron detection probability region has no clear boundaries. However, it is possible to single out the space where the probability of finding an electron is maximum.

The space around the atomic nucleus, in which the electron is most likely to be found, is called the 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. According to the form, $4$ of currently known types of orbitals are distinguished, which are denoted by the Latin letters $s, p, d$ and $f$. A graphic representation of some forms of electronic orbitals is shown in the figure.

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

An 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. The electrons of the first energy level, closest to the nucleus, have the lowest energy. Compared with the electrons of the first level, the electrons of the next levels are characterized by a large amount of energy. Consequently, the electrons of the outer level are the least strongly bound to the nucleus of the atom.

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

The largest number of electrons in the 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: the first energy level closest to the nucleus can contain 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 the energy levels (electronic layers) arranged?

Starting from the second energy level $(n = 2)$, each of the levels is subdivided into sublevels (sublayers), slightly different from each other by 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; the fourth is four. Sublevels, in turn, are formed by orbitals.

Each value of $n$ corresponds to the number of orbitals equal to $n^2$. According to the data presented in the table, it is possible to trace the relationship between the principal quantum number $n$ and the number of sublevels, the type and number of orbitals, and the maximum number of electrons per sublevel and level.

Principal quantum number, types and number of orbitals, maximum number of electrons at sublevels and levels.

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

It is customary to designate sublevels in 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 for 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.

atom nucleus

But not only electrons are part of atoms. Physicist Henri Becquerel discovered that a natural mineral containing uranium salt also emits unknown radiation, illuminating photographic films that are closed from light. This phenomenon has been 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 are a stream of electrons;
3. $γ$-rays are electromagnetic waves with a negligible mass that do not carry an electric charge.

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

How is the atom arranged?

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

By placing the 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, that area in which there are electrons that have a negative charge. If we apply a figurative comparison, then the entire volume of the atom can be likened to the Luzhniki stadium, 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, such a 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 electrons, but opposite in sign $(+1)$, and a mass equal to the mass of a hydrogen atom (it is accepted in chemistry as a unit). Protons are denoted by $↙(1)↖(1)p$ (or $р+$). 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 denoted by $↙(0)↖(1)n$ (or $n^0$).

Protons and neutrons are collectively 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:

Since the mass of the electron, which is negligible, can be neglected, it is obvious that the entire mass of the atom is concentrated in the nucleus. Electrons are denoted 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. And how to determine the number of neutrons?

As you know, the mass of an atom is the sum of the mass of protons and neutrons. Knowing the ordinal 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:

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

$56 – 26 = 30$.

The table shows 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- the same and topos- place, means "occupying one place" (cell) in the Periodic system of elements.

Chemical elements found in nature are a mixture of isotopes. Thus, carbon has three isotopes with a mass of $12, 13, 14$; oxygen - three isotopes with a mass of $16, 17, 18$, etc.

Usually given in the Periodic system, the relative atomic mass of a chemical element 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 $37$ (there are $25%$); 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 the isotopes of most chemical elements, such as potassium, argon:

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

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

Now it is possible to 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 electron shells of atoms of the elements of the first four periods

Consider the mapping of the electronic configurations of the atoms of the elements by the periods of the system of D. I. Mendeleev.

Elements of the first period.

Schemes of the electronic structure of atoms show the distribution of electrons over electronic layers (energy levels).

The electronic formulas of atoms show the distribution of electrons over energy levels and sublevels.

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

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

Hydrogen and helium are $s$-elements, these atoms have $s$-orbitals filled with electrons.

Elements of the second period.

For all elements of the second period, the first electron layer is filled, and the 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 rules of Pauli and Hund.

In the neon atom, the second electron layer is complete - it has $8$ 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 3s-, 3p- and 3d-sublevels.

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

A $3.5$-electron orbital is completed at the magnesium atom. $Na$ and $Mg$ are $s$-elements.

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

$↙(18)(Ar)$ Argon

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

In an argon atom, the outer layer (the third electron layer) has $8$ electrons. 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 $Ar$ - $p$ -elements.

$s-$ and $r$ -elements form main subgroups in the Periodic system.

Elements of the fourth period.

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

1. we denote conditionally the graphic electronic formula of argon as follows: $Ar$;
2. we will not depict the sublevels that are not filled for 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 pre-external electron layer is filled, they are referred to transition elements.

Pay attention to the structure of the electron shells of chromium and copper atoms. In them, one electron "falls" from the $4s-$ to the $3d$ sublevel, which is explained by the greater energy stability of the resulting $3d^5$ and $3d^(10)$ electronic configurations:

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

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

Element symbol, serial number, name	Diagram of the electronic structure	Electronic formula	Graphic 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 $1s^2(2)s^2(2)p ^6(3)p^6(3)d^5(4)s^1$
$↙(29)(Сu)$ Chromium		$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^(10)$ 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 the $3s, 3p$ and $3d$ sublevels are filled in it, in total there are $18$ of electrons on them.

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

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

The elements of the fifth period are filling the sublevels in the following order: $5s → 4d → 5р$. And there are also exceptions related to the "failure" of electrons, for $↙(41)Nb$, $↙(42)Mo$, $↙(44)Ru$, $↙(45)Rh$, $↙(46) Pd$, $↙(47)Ag$. $f$ appear in the sixth and seventh periods -elements, i.e. elements whose $4f-$ and $5f$-sublevels of the third outside electronic layer are being filled, respectively.

$4f$ -elements called lanthanides.

$5f$ -elements called actinides.

The order of filling of electronic sublevels in the atoms of elements of the sixth period: $↙(55)Cs$ and $↙(56)Ba$ - $6s$-elements; $↙(57)La ... 6s^(2)5d^(1)$ - $5d$-element; $↙(58)Ce$ – $↙(71)Lu - 4f$-elements; $↙(72)Hf$ – $↙(80)Hg - 5d$-elements; $↙(81)Т1$ – $↙(86)Rn - 6d$-elements. But even here there are elements in which the order of filling of electron 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 electronic 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. $r$ -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 preexternal level of the atom is filled with electrons; $d$-elements include elements of secondary subgroups of groups I–VIII, i.e. elements of intercalated decades of large periods located between $s-$ and $p-$elements. They are also called transition elements;
4. $f$ -elements;$f-$sublevel of the third level of the atom outside is filled with electrons; these include lanthanides and actinides.

The electronic configuration of the atom. Ground and excited states of atoms

The Swiss physicist W. Pauli in $1925$ established that An atom can have at most two electrons in one orbital. having opposite (antiparallel) spins (translated from English as a spindle), i.e. possessing such properties that can be conditionally imagined 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 an orbital, then it is called unpaired, if two, then this paired electrons, i.e. electrons with opposite spins.

The figure shows a diagram of the division of energy levels into sublevels.

$s-$ Orbital, as you already know, has a spherical shape. The hydrogen atom electron $(n = 1)$ is located on this orbital and is unpaired. According to this his 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 (orbital type), and the number that is written to the 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 the same $s-$orbital, this formula is: $1s^2$. The electron shell of the helium atom is complete and very stable. Helium is a noble gas. The second energy level $(n = 2)$ has four orbitals, one $s$ and three $p$. Second-level $s$-orbital electrons ($2s$-orbitals) have a 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 amount of electron energy on it and, therefore, with a corresponding diameter, growing as the value of $n$.$s-$Orbital increases, as you already know , has a spherical shape. The hydrogen atom electron $(n = 1)$ is located on 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 (orbital type), and the number that is written to the 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 the same $s-$orbital, this formula is: $1s^2$. The electron shell of the helium atom is complete and very stable. Helium is a noble gas. The second energy level $(n = 2)$ has four orbitals, one $s$ and three $p$. Electrons of $s-$orbitals of the second level ($2s$-orbitals) have a 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 amount of electron energy on it and, therefore, with a corresponding diameter, growing as the value of $n$ increases.

$r-$ Orbital It has the shape of a dumbbell, or volume eight. All three $p$-orbitals are located in the atom mutually perpendicularly along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized again that each energy level (electronic layer), starting from $n= 2$, has three $p$-orbitals. As the value of $n$ increases, the 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 weaker bound to the atomic nucleus, so a lithium atom can easily give it away (as you probably remember, this process is called oxidation), turning into a lithium ion $Li^+$.

In the beryllium atom Be, the fourth electron is also placed 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.

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

For elements of the third period, $3s-$ 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 above full electronic formulas, 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 respectively $4s-$ and $5s$-orbitals: $↙(19)K 2, 8, 8, 1;$ $↙(38)Sr 2, 8, 18, 8, 2$. Starting from the third element of each large period, the next ten electrons will go to the previous $3d-$ and $4d-$orbitals, respectively (for elements of secondary 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 (respectively $4p-$ and $5p-$) $p-$sublevel will start 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, as follows: 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 $14$ of electrons will enter the third energy level from the outside, the $4f$ and $5f$ orbitals of the lantonides and actinides, respectively: $↙(64)Gd 2, 8, 18, 25, 9, 2;$ $↙(92 )U 2, 8, 18, 32, 21, 9, 2$.

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

Very often, the structure of the electron shells of atoms is depicted using energy or quantum cells - they write down the so-called graphic electronic formulas. For this record, the following notation is used: each quantum cell is denoted by a cell that corresponds to one orbital; each electron is indicated by an arrow corresponding to the direction of the spin. When writing a graphical electronic formula, two rules should be remembered: Pauli principle, according to which a cell (orbital) can have no more than two electrons, 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 already be oppositely directed.

(Lecture notes)

The structure of the atom. Introduction.

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

1. The structure of the atom and the Periodic system of elements

DI. Mendeleev.

Knowledge and understanding of the regularities of the order of filling the Periodic system of elements D.I. Mendeleev allows us to understand the following:

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

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

3. the ability and "ease" of an element to give or receive electrons when interacting with another element,

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

The structure of the atom.

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

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

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

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

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

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

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

Fig.1. 1s orbital

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

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

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

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

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

When filling three p-AOs, five d-AOs and seven f-AOs with electrons, one should be guided not only by the Pauli principle but also by the Hund rule: The filling of the orbitals of one subshell in the ground state occurs with electrons with the same spins.

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

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

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

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

Currently, 109 elements are known.

2. Ionization energy, electron affinity and electronegativity.

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

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

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

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

Table 1

The first ionization energy (EI), electron affinity (SE) and electronegativity χ) of atoms in the periodic system.

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

χ - Pauling electronegativity

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

The lesson is devoted to the formation of ideas about the complex structure of the atom. The state of electrons in an atom is considered, the concepts of "atomic orbital and electron cloud", the forms of orbitals (s--, p-, d-orbitals) are introduced. Also considered are aspects such as the maximum number of electrons at energy levels and sublevels, the distribution of electrons over energy levels and sublevels in atoms of elements of the first four periods, valence electrons of s-, p- and d-elements. A graphical diagram of the structure of the electronic layers of atoms (electron-graphic formula) is given.

Topic: The structure of the atom. Periodic law D.I. Mendeleev

Lesson: The structure of the atom

Translated from Greek, the word " atom" means "indivisible". However, phenomena have been discovered that demonstrate the possibility of its division. These are the emission of x-rays, the emission of cathode rays, the phenomenon of the photoelectric effect, the phenomenon of radioactivity. Electrons, protons, and neutrons are the particles that make up an atom. They're called subatomic particles.

Tab. one

In addition to protons, the nucleus of most atoms contains neutrons that carry no charge. As can be seen from Table. 1, the mass of the neutron practically does not differ from the mass of the proton. Protons and neutrons make up the nucleus of an atom and are called nucleons (nucleus - nucleus). Their charges and masses in atomic mass units (a.m.u.) are shown in Table 1. When calculating the mass of an atom, the mass of an electron can be neglected.

Mass of an atom ( mass number) is equal to the sum of the masses of the protons and neutrons that make up its nucleus. The mass number is denoted by the letter BUT. From the name of this quantity, it can be seen that it is closely related to the atomic mass of the element rounded to an integer. A=Z+N

Here A- mass number of an atom (the sum of protons and neutrons), Z- nuclear charge (number of protons in the nucleus), N is the number of neutrons in the nucleus. According to the doctrine of isotopes, the concept of "chemical element" can be given the following definition:

chemical element A group of atoms with the same nuclear charge is called.

Some elements exist as multiple isotopes. "Isotopes" means "occupying the same place." Isotopes have the same number of protons, but differ in mass, i.e., the number of neutrons in the nucleus (number N). Because neutrons have little to no effect on the chemical properties of elements, all isotopes of the same element are chemically indistinguishable.

Isotopes are called varieties of atoms of the same chemical element with the same nuclear charge (that is, with the same number of protons), but with a different number of neutrons in the nucleus.

Isotopes differ from each other only in mass number. This is indicated either by a superscript in the right corner, or in a line: 12 C or C-12 . If an element contains several natural isotopes, then in the periodic table D.I. Mendeleev indicates its average atomic mass, taking into account the prevalence. For example, chlorine contains 2 natural isotopes 35 Cl and 37 Cl, the content of which is 75% and 25%, respectively. Thus, the atomic mass of chlorine will be equal to:

BUTr(Cl)=0,75 . 35+0,25 . 37=35,5

For artificially synthesized heavy atoms, one atomic mass value is given in square brackets. This is the atomic mass of the most stable isotope of that element.

Basic models of the structure of the atom

Historically, the Thomson model of the atom was the first in 1897.

Rice. 1. Model of the structure of the atom by J. Thomson

The English physicist J. J. Thomson suggested that atoms consist of a positively charged sphere in which electrons are interspersed (Fig. 1). This model is figuratively called "plum pudding", a bun with raisins (where "raisins" are electrons), or "watermelon" with "seeds" - electrons. However, this model was abandoned, since experimental data were obtained that contradicted it.

Rice. 2. Model of the structure of the atom by E. Rutherford

In 1910, the English physicist Ernst Rutherford, with his students Geiger and Marsden, conducted an experiment that gave amazing results that were inexplicable from the point of view of the Thomson model. Ernst Rutherford proved by experience that in the center of the atom there is a positively charged nucleus (Fig. 2), around which, like planets around the Sun, electrons revolve. The atom as a whole is electrically neutral, and the electrons are held in the atom due to the forces of electrostatic attraction (Coulomb forces). This model had many contradictions and, most importantly, did not explain why electrons do not fall on the nucleus, as well as the possibility of absorption and emission of energy by it.

The Danish physicist N. Bohr in 1913, taking Rutherford's model of the atom as a basis, proposed a model of the atom in which electron-particles revolve around the atomic nucleus in much the same way as the planets revolve around the Sun.

Rice. 3. Planetary model of N. Bohr

Bohr suggested that electrons in an atom can only exist stably in orbits at strictly defined distances from the nucleus. These orbits he called stationary. An electron cannot exist outside stationary orbits. Why this is so, Bohr could not explain at the time. But he showed that such a model (Fig. 3) makes it possible to explain many experimental facts.

Currently used to describe the structure of the atom quantum mechanics. This is a science, the main aspect of which is that the electron has the properties of a particle and a wave at the same time, i.e., wave-particle duality. According to quantum mechanics, the region of space in which the probability of finding an electron is greatest is calledorbital. The farther the electron is from the nucleus, the lower its interaction energy with the nucleus. Electrons with similar energies form energy level. Number of energy levels equals period number, in which this element is located in the table D.I. Mendeleev. There are various shapes of atomic orbitals. (Fig. 4). The d-orbital and f-orbital have a more complex shape.

Rice. 4. Shapes of atomic orbitals

There are exactly as many electrons in the electron shell of any atom as there are protons in its nucleus, so the atom as a whole is electrically neutral. Electrons in an atom are arranged so that their energy is minimal. The farther the electron is from the nucleus, the more orbitals and the more complex they are in shape. Each level and sublevel can only hold a certain number of electrons. The sublevels, in turn, consist of orbitals.

At the first energy level, closest to the nucleus, there can be one spherical orbital ( 1 s). At the second energy level - a spherical orbital, large in size and three p-orbitals: 2 s2 ppp. On the third level: 3 s3 ppp3 dddd.

In addition to movement around the nucleus, electrons also have movement, which can be represented as their movement around their own axis. This rotation is called spin ( in lane from English. "spindle"). Only two electrons with opposite (antiparallel) spins can be in one orbital.

Maximum number of electrons per energy level is determined by the formula N=2 n 2.

Where n is the main quantum number (energy level number). See table. 2

Tab. 2

Depending on which orbital the last electron is in, they distinguish s-, p-, d-elements. Elements of the main subgroups belong to s-, p-elements. In the side subgroups are d-elements

Graphic diagram of the structure of the electronic layers of atoms (electronic graphic formula).

To describe the arrangement of electrons in atomic orbitals, the electronic configuration is used. To write it in a line, orbitals are written in the legend ( s--, p-, d-,f-orbitals), and in front of them are numbers indicating the number of the energy level. The larger the number, the further the electron is from the nucleus. In upper case, above the designation of the orbital, the number of electrons in this orbital is written (Fig. 5).

Rice. 5

Graphically, the distribution of electrons in atomic orbitals can be represented as cells. Each cell corresponds to one orbital. There will be three such cells for the p-orbital, five for the d-orbital, and seven for the f-orbital. One cell can contain 1 or 2 electrons. According to Gund's rule, electrons are distributed in orbitals of the same energy (for example, in three p-orbitals), first one at a time, and only when there is already one electron in each such orbital, the filling of these orbitals with second electrons begins. Such electrons are called paired. This is explained by the fact that in neighboring cells, electrons repel each other less, as similarly charged particles.

See fig. 6 for atom 7 N.

Rice. 6

The electronic configuration of the scandium atom

21 sc: 1 s 2 2 s 2 2 p 6 3 s 2 3 p 6 4 s 2 3 d 1

Electrons in the outer energy level are called valence electrons. 21 sc refers to d-elements.

Summing up the lesson

At the lesson, the structure of the atom, the state of electrons in the atom were considered, the concept of "atomic orbital and electron cloud" was introduced. Students learned what the shape of orbitals is ( s-, p-, d-orbitals), what is the maximum number of electrons at energy levels and sublevels, the distribution of electrons over energy levels, what is s-, p- and d-elements. A graphical diagram of the structure of the electronic layers of atoms (electron-graphic formula) is given.

Bibliography

1. Rudzitis G.E. Chemistry. Fundamentals of General Chemistry. Grade 11: textbook for educational institutions: basic level / G.E. Rudzitis, F.G. Feldman. - 14th ed. - M.: Education, 2012.

2. Popel P.P. Chemistry: 8th grade: a textbook for general educational institutions / P.P. Popel, L.S. Krivlya. - K .: Information Center "Academy", 2008. - 240 p.: ill.

3. A.V. Manuilov, V.I. Rodionov. Fundamentals of chemistry. Internet tutorial.

Homework

1. No. 5-7 (p. 22) Rudzitis G.E. Chemistry. Fundamentals of General Chemistry. Grade 11: textbook for educational institutions: basic level / G.E. Rudzitis, F.G. Feldman. - 14th ed. - M.: Education, 2012.

2. Write electronic formulas for the following elements: 6 C, 12 Mg, 16 S, 21 Sc.

3. Elements have the following electronic formulas: a) 1s 2 2s 2 2p 4 .b) 1s 2 2s 2 2p 6 3s 2 3p 1. c) 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2 . What are these elements?

An atom is the smallest particle of matter. Its study began in ancient Greece, when the attention of not only scientists, but also philosophers was riveted to the structure of the atom. What is the electronic structure of an atom, and what basic information is known about this particle?

The structure of the atom

Already ancient Greek scientists guessed the existence of the smallest chemical particles that make up any object and organism. And if in the XVII-XVIII centuries. chemists were sure that the atom is an indivisible elementary particle, then at the turn of the 19th-20th centuries, they managed to prove experimentally that the atom is not indivisible.

An atom, being a microscopic particle of matter, consists of a nucleus and electrons. The nucleus is 10,000 times smaller than an atom, but almost all of its mass is concentrated in the nucleus. The main characteristic of the atomic nucleus is that it has a positive charge and is made up of protons and neutrons. Protons are positively charged, while neutrons have no charge (they are neutral).

They are connected to each other by the strong nuclear force. The mass of a proton is approximately equal to the mass of a neutron, but at the same time it is 1840 times greater than the mass of an electron. Protons and neutrons have a common name in chemistry - nucleons. The atom itself is electrically neutral.

An atom of any element can be denoted by an electronic formula and an electronic graphic formula:

Rice. 1. Electron-graphic formula of the atom.

The only chemical element from the periodic table, the nucleus of which does not contain neutrons, is light hydrogen (protium).

An electron is a negatively charged particle. The electron shell consists of electrons moving around the nucleus. Electrons have properties to be attracted to the nucleus, and between each other they are influenced by the Coulomb interaction. To overcome the attraction of the nucleus, the electrons must receive energy from an external source. The farther the electron is from the nucleus, the less energy is needed for this.

Atom Models

For a long time, scientists have sought to understand the nature of the atom. At an early stage, the ancient Greek philosopher Democritus made a great contribution. Although now his theory seems banal and too simple to us, at a time when the concept of elementary particles was just beginning to emerge, his theory of pieces of matter was taken quite seriously. Democritus believed that the properties of any substance depend on the shape, mass and other characteristics of atoms. So, for example, near fire, he believed, there are sharp atoms - therefore, fire burns; water has smooth atoms, so it can flow; in solid objects, in his view, the atoms were rough.

Democritus believed that absolutely everything consists of atoms, even the human soul.

In 1904, J. J. Thomson proposed his model of the atom. The main provisions of the theory boiled down to the fact that the atom was represented as a positively charged body, inside of which there were electrons with a negative charge. Later this theory was refuted by E. Rutherford.

Rice. 2. Thomson's model of the atom.

Also in 1904, the Japanese physicist H. Nagaoka proposed an early planetary model of the atom by analogy with the planet Saturn. According to this theory, electrons are united in rings and revolve around a positively charged nucleus. This theory turned out to be wrong.

In 1911, E. Rutherford, having done a series of experiments, concluded that the atom in its structure is similar to the planetary system. After all, electrons, like planets, move in orbits around a heavy positively charged nucleus. However, this description contradicted classical electrodynamics. Then the Danish physicist Niels Bohr in 1913 introduced the postulates, the essence of which was that the electron, being in some special states, does not radiate energy. Thus, Bohr's postulates showed that classical mechanics is inapplicable to atoms. The planetary model described by Rutherford and supplemented by Bohr was called the Bohr-Rutherford planetary model.

Rice. 3. Bohr-Rutherford planetary model.

Further study of the atom led to the creation of such a section as quantum mechanics, with the help of which many scientific facts were explained. Modern ideas about the atom have developed from the Bohr-Rutherford planetary model. Evaluation of the report

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