Electronic structure of the atom. Electronic configuration of an atom
The concept of atom arose in the ancient world to denote particles of matter. Translated from Greek, atom means “indivisible.”
Electrons
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 the current source with two metal plates - the cathode and the anode, soldered into glass tube, from which the air was pumped out. When a voltage of about 10 thousand volts was applied to the electrode plates, a luminous discharge flashed in the tube, and from the cathode (negative pole) to the anode ( positive pole) particles were flying, 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 core, and the totality various provisions it 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 is $≈30%$, etc. There is some uncertainty in the state of the electron. To characterize this special condition, 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.
Space around atomic nucleus, in which the 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 $f$. Graphic image Some forms of electron orbitals are shown in the figure.
The most important characteristic of the motion of an electron in a certain orbital is 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 that are protected from light. This phenomenon was called radioactivity.
There are three types of radioactive rays:
- $α$-rays, which consist of $α$-particles having a charge $2$ times greater than the charge of an electron, but with positive sign, and the mass is $4$ times greater than the mass of the hydrogen atom;
- $β$-rays represent a flow of electrons;
- $γ$-rays - electromagnetic waves with negligible mass, not carrying an electrical charge.
Therefore, the atom has complex structure- consists of a positively charged nucleus and electrons.
How is an 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 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, the region in which electrons with a negative charge are located. If you apply figurative comparison, then the entire volume of the atom can be likened to the stadium in Luzhniki, and the nucleus 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 two types of particles - protons and neutrons.
Protons have a charge equal to the charge of electrons, but opposite in sign $(+1)$, and mass, equal to mass hydrogen atom (it is taken as a unit 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 shows the main characteristics 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 $37$ (there 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 differ greatly in properties due to a sharp multiple increase in their relative atomic mass; they were even given 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 strict and scientific definition 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.
Schemes electronic structure 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 unfilled $3d$-orbitals.
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:
- Let us denote the conventional graphical electronic formula of argon as follows: $Ar$;
- 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 $3d^5$ 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 | 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.
The elements of the fifth period is running filling in sublevels in the following order: $5s → 4d → 5р$. 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:
- $s$ -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;
- $r$ -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;
- $d$ -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;
- $f$ -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 before the letter $(1...)$, Latin letter denote a sublevel (type of orbital), 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 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 the 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 at the same time have same value back, and only then mate, but the backs, according to the Pauli principle, will already be in opposite directions.
(Lecture notes)
The structure of the atom. Introduction.
The object of study in chemistry is chemical elements and their compounds. Chemical element called a collection of atoms with the same positive charge. Atom- This smallest particle chemical element, preserving it chemical properties. By bonding with each other, atoms of the same or different elements form more complex particles - molecules. A collection of atoms or molecules form chemical substances. Each individual chemical substance is characterized by a set of individual physical properties, such as boiling and melting points, density, electrical and thermal conductivity, etc.
1. Atomic structure and the Periodic Table of Elements
DI. Mendeleev.
Knowledge and understanding of the laws of the order of filling the Periodic Table of Elements D.I. Mendeleev allows us to understand the following:
1. the physical essence of the existence of certain elements in nature,
2. the nature of the chemical valence of the element,
3. the ability and “lightness” of an element to give or accept electrons when interacting with another element,
4. the nature of the chemical bonds that a given element can form when interacting with other elements, the spatial structure of simple and complex molecules, etc., etc.
The structure of the atom.
An atom is a complex microsystem of elementary particles in motion and interacting with each other.
In the late 19th and early 20th centuries, it was discovered that atoms are made up of smaller particles: neutrons, protons and electrons. The last two particles are charged particles, the proton carries a positive charge, the electron a negative one. Since the atoms of an element in the ground state are electrically neutral, this means that the number of protons in an atom of any element is equal to the number of electrons. The mass of atoms is determined by the sum of the masses of protons and neutrons, the number of which is equal to the difference between the mass of the atoms and its serial number in the periodic system D.I. Mendeleev.
In 1926, Schrödinger proposed describing the movement of microparticles in the atom of an element using the wave equation he derived. When solving the Schrödinger wave equation for the hydrogen atom, three integer quantum numbers appear: n, ℓ And m ℓ , which characterize the state of the electron in three-dimensional space in the central field of the nucleus. Quantum numbers n, ℓ And m ℓ take integer values. Wave function defined by three quantum numbers n, ℓ And m ℓ and obtained as a result of solving the Schrödinger equation is called an orbital. An orbital is a region of space in which an electron is most likely to be found, belonging to an atom of a chemical element. Thus, solving the Schrödinger equation for the hydrogen atom leads to the appearance of three quantum numbers, physical meaning which is that they characterize three different types of orbitals that an atom can have. Let's take a closer look at each quantum number.
Principal quantum number n can take any positive integer values: n = 1,2,3,4,5,6,7...It characterizes the energy of the electron level and the size of the electron “cloud”. It is characteristic that the number of the main quantum number coincides with the number of the period in which the element is located.
Azimuthal or orbital quantum numberℓ can take integer values from ℓ = 0….to n – 1 and determines the moment of electron motion, i.e. orbital shape. For various numerical valuesℓ use the following notation: ℓ = 0, 1, 2, 3, and are indicated by the symbols s, p, d, f, respectively for ℓ = 0, 1, 2 and 3. In the periodic table of elements there are no elements with a spin number ℓ = 4.
Magnetic quantum numberm ℓ characterizes the spatial arrangement of electron orbitals and, consequently, the electromagnetic properties of the electron. It can take values from – ℓ to + ℓ , including zero.
The shape, or more precisely, the symmetry properties of atomic orbitals depend on quantum numbers ℓ And m ℓ . "Electronic cloud" corresponding s- the orbitals have, have the shape of a ball (at the same time ℓ = 0).
Fig.1. 1s orbital
The orbitals defined by the quantum numbers ℓ = 1 and m ℓ = -1, 0 and +1 are called p-orbitals. Since m ℓ in this case has three different meanings, then the atom has three energetically equivalent p-orbitals (the 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 look like three-dimensional figure eights, oriented along the x, y and z axes in an external field (Fig. 1.2). Hence the origin of the symbolism p x , p y and p z .
Fig.2. p x, p y and p z orbitals
In addition, there are d- and f- atomic orbitals, for the first ℓ = 2 and m ℓ = -2, -1, 0, +1 and +2, i.e. five AOs, for the second ones ℓ = 3 and m ℓ = -3, -2, -1, 0, +1, +2 and +3, i.e. 7 JSC.
Fourth quantum m s called the spin quantum number, was introduced to explain certain subtle effects in the spectrum of the hydrogen atom by Goudsmit and Uhlenbeck in 1925. The spin of an electron is the angular momentum of a charged elementary particle of an electron, the orientation of which is quantized, i.e. strictly limited to certain angles. This orientation is determined by the value of the spin magnetic quantum number (s), which for the electron is equal to ½ , therefore for the electron according to the quantization rules m s = ± ½. In this regard, to the set of three quantum numbers we should add the quantum number m s . Let us emphasize once again that four quantum numbers determine the order of construction of Mendeleev’s periodic table of elements and explain why there are only two elements in the first period, eight in the second and third, 18 in the fourth, etc. However, in order to explain the structure of many-electron atoms, the order of filling electronic levels as the positive charge of the atom increases, it is not enough to have an idea of the four quantum numbers that “control” the behavior of electrons when filling electron orbitals, but you need to know some more simple rules, namely, Pauli's principle, Hund's rule and Kleczkowski's rules.
According to the Pauli principle In the same quantum state, characterized by certain values of four quantum numbers, there cannot be more than one electron. This means that one electron can, in principle, be placed in any atomic orbital. Two electrons can be in the same atomic orbital only if they have different spin quantum numbers.
When filling three p-AOs, five d-AOs and seven f-AOs with electrons, one should be guided, in addition to the Pauli principle, by Hund’s rule: The filling of the orbitals of one subshell in the ground state occurs with electrons with identical spins.
When filling the subshells (p, d, f)the absolute value of the sum of spins must be maximum.
Klechkovsky's rule. According to Klechkovsky’s rule, when fillingd And felectron orbital must be respectedprinciple of minimum energy. According to this principle, electrons in the ground state occupy orbitals with minimal energy levels. The energy of a sublevel is determined by the sum of quantum numbersn + ℓ = E .
Klechkovsky's first rule: First, those sublevels for whichn + ℓ = E minimal.
Klechkovsky's second rule: in case of equalityn + ℓ for several sublevels, the sublevel for which is filledn minimal .
Currently, 109 elements are known.
2. Ionization energy, electron affinity and electronegativity.
The most important characteristics of the electronic configuration of an atom are ionization energy (IE) or ionization potential (IP) and the atom's electron affinity (EA). Ionization energy is the change in energy during the removal of an electron from a free atom at 0 K: A = + + ē . The dependence of ionization energy on the atomic number Z of an element and the size of the atomic radius has a pronounced periodic character.
Electron affinity (EA) is the change in energy that accompanies the addition of an electron to an isolated atom to form a negative ion at 0 K: A + ē = A - (the atom and ion are in their ground states). In this case, the electron occupies the lowest vacant atomic orbital (LUAO) if the VZAO is occupied by two electrons. The SE strongly depends on their orbital electronic configuration.
Changes in EI and SE correlate with changes in many properties of elements and their compounds, which is used to predict these properties from EI and SE values. The 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 electronic layers and which correlates well with an increase in the reducing power of the element.
Table 1 of the Periodic Table of Elements shows the values of EI and SE in eV/per atom. Note that exact values SEs are known only for a few atoms; their values are highlighted in Table 1.
Table 1
First ionization energy (EI), electron affinity (EA) and electronegativity χ) of atoms in the periodic table.
|
χ |
0.747 2. 1 0 0, 3 7 |
1,2 2 |
||||||||||||||||
|
χ |
0.54 1. 55 |
-0.3 1. 1 3 |
0.2 0. 91 |
1.2 5 |
-0. 1 0, 55 |
1.47 0. 59 |
3.45 0. 64 |
1 ,60 |
||||||||||
|
χ |
0. 7 4 1. 89 |
-0.3 1 . 3 1 1 . 6 0 |
0. 6 |
1.63 |
0.7 |
2.07 |
3.61 | |||||||||||
|
χ |
2.3 6 |
- 0 .6 |
1.26(α) |
-0.9 1 . 39 |
0. 18 |
1.2 |
0. 6 |
2.07 |
3.36 | |||||||||
|
χ |
2.4 8 |
-0.6 1 . 56 |
0. 2 |
2.2 | ||||||||||||||
|
χ |
2.6 7 |
2, 2 1 |
ABOUTs |
χ – electronegativity according to Pauling
r- atomic radius, (from “Laboratory and seminar classes in general and inorganic chemistry”, N.S. Akhmetov, M.K. Azizova, L.I. Badygina)
The concept of “atom” has been familiar to humanity since the times of Ancient Greece. According to the statement of ancient philosophers, an atom is the smallest particle that is part of a substance.
Electronic structure of the atom
An atom consists of a positively charged nucleus containing protons and neutrons. Electrons move in orbits around the nucleus, each of which can be characterized by a set of four quantum numbers: principal (n), orbital (l), magnetic (ml) and spin (ms or s).
The principal quantum number determines the energy of the electron and the size of the electron clouds. The energy of an electron mainly depends on the distance of the electron from the nucleus: the closer the electron is to the nucleus, the lower its energy. In other words, the principal quantum number determines the location of the electron at a particular energy level (quantum layer). The principal quantum number has the values of a series of integers from 1 to infinity.
The orbital quantum number characterizes the shape of the electron cloud. Various shape electron clouds causes a change in the energy of electrons within one energy level, i.e. splitting it into energy sublevels. The orbital quantum number can have values from zero to (n-1), for a total of n values. Energy sublevels are designated by letters:
The magnetic quantum number shows the orientation of the orbital in space. It accepts any integer numeric value from (+l) to (-l), including zero. Number possible values magnetic quantum number is equal to (2l+1).
An electron, moving in the field of the atomic nucleus, in addition to the orbital angular momentum, also has its own angular momentum, which characterizes its spindle-shaped rotation around its own axis. This property of an electron is called spin. The magnitude and orientation of the spin is characterized by the spin quantum number, which can take values (+1/2) and (-1/2). Positive and negative values the back is related to its direction.
Before all of the above became known and confirmed experimentally, there were several models of the structure of the atom. One of the first models of the structure of the atom was proposed by E. Rutherford, who, in experiments on the scattering of alpha particles, showed that almost the entire mass of the atom is concentrated in a very small volume - a positively charged nucleus. According to his model, electrons move around the nucleus at a sufficiently large distance, and their number is such that, on the whole, the atom is electrically neutral.
Rutherford's model of the structure of the atom was developed by N. Bohr, who in his research also combined Einstein's teachings on light quanta and quantum theory Planck radiation. Finished what we started and presented it to the world modern model structure of the atom of a chemical element Louis de Broglie and Schrödinger.
Examples of problem solving
EXAMPLE 1
| Exercise | List the number of protons and neutrons contained in the nuclei of nitrogen (atomic number 14), silicon (atomic number 28), and barium (atomic number 137). |
| Solution | The number of protons in the nucleus of an atom of a chemical element is determined by its serial number in the Periodic Table, and the number of neutrons is the difference between the mass number (M) and the charge of the nucleus (Z). Nitrogen: n(N)= M -Z = 14-7 = 7. Silicon: n(Si)= M -Z = 28-14 = 14. Barium: n (Ba)= M -Z = 137-56 = 81. |
| Answer | The number of protons in the nitrogen nucleus is 7, neutrons - 7; in the nucleus of a silicon atom there are 14 protons and 14 neutrons; In the nucleus of a barium atom there are 56 protons and 81 neutrons. |
EXAMPLE 2
| Exercise | Arrange the energy sublevels in the order in which they are filled with electrons: a) 3p, 3d, 4s, 4p; b) 4d , 5s, 5p, 6s; c) 4f , 5s , 6r; 4d , 6s; d) 5d, 6s, 6p, 7s, 4f . |
|
| Solution | Energy sublevels are filled with electrons in accordance with Klechkovsky's rules. Required condition is the minimum value of the sum of the principal and orbital quantum numbers. The s-sublevel is characterized by the number 0, p - 1, d - 2 and f-3. The second condition is that the sublevel with the smallest value of the principal quantum number is filled first. | |
| Answer | a) Orbitals 3p, 3d, 4s, 4p will correspond to the numbers 4, 5, 4 and 5. Consequently, filling with electrons will occur in the following sequence: 3p, 4s, 3d, 4p. b) 4d orbitals , 5s, 5p, 6s will correspond to the numbers 7, 5, 6 and 6. Therefore, filling with electrons will occur in the following sequence: 5s, 5p, 6s, 4d. c) Orbitals 4f , 5s , 6r; 4d , 6s will correspond to the numbers 7, 5, 76 and 6. Therefore, filling with electrons will occur in the following sequence: 5s, 4d , 6s, 4f, 6r. d) Orbitals 5d, 6s, 6p, 7s, 4f will correspond to the numbers 7, 6, 7, 7 and 7. Consequently, filling with electrons will occur in the following sequence: 6s, 4f, 5d, 6p, 7s. |
Everything in the world is made of atoms. But where did they come from, and what are they made of? Today we answer these simple and fundamental questions. After all, many people living on the planet say that they do not understand the structure of the atoms from which they themselves are composed.
Naturally, dear reader understands that in this article we try to present everything at the simplest and most interesting level, so we do not “load” it with scientific terms. For those who want to study the issue in more detail professional level, we recommend reading specialized literature. Nevertheless, the information in this article can serve well in your studies and simply make you more erudite.
An atom is a particle of a substance of microscopic size and mass, the smallest part of a chemical element, which is the carrier of its properties. In other words, this smallest particle a substance that can enter into chemical reactions.
Discovery history and structure
The concept of an atom was known back in Ancient Greece. Atomism – physical theory, which states that all material objects consist of indivisible particles. Along with Ancient Greece, the ideas of atomism also developed in parallel in Ancient India.
It is not known whether the aliens told the philosophers of that time about atoms, or whether they came up with it themselves, but chemists were able to experimentally confirm this theory much later - only in the seventeenth century, when Europe emerged from the abyss of the Inquisition and the Middle Ages.
For a long time, the dominant idea of the structure of the atom was the idea of it as an indivisible particle. The fact that the atom can still be divided became clear only at the beginning of the twentieth century. Rutherford, thanks to his famous experiment with the deflection of alpha particles, learned that the atom consists of a nucleus around which electrons revolve. The planetary model of the atom was adopted, according to which electrons rotate around the nucleus, like our planets solar system around the star.
Modern representations much progress has been made about the structure of the atom. The nucleus of an atom, in turn, consists of subatomic particles, or nucleons - protons and neutrons. It is nucleons that make up the bulk of the atom. Moreover, protons and neutrons are also not indivisible particles, and consist of fundamental particles - quarks.
The nucleus of an atom has a positive electric charge, and electrons rotating in orbit are negative. Thus, the atom is electrically neutral.
Below we give an elementary diagram of the structure of the carbon atom.
Properties of atoms
Weight
The mass of atoms is usually measured in atomic mass units - a.m.u. Atomic unit mass is the mass of 1/12 of the freely resting carbon atom in the ground state.
In chemistry, the concept is used to measure the mass of atoms "moth". 1 mole is the amount of substance that contains the number of atoms equal to the number Avogadro.
Size
The sizes of atoms are extremely small. So, the smallest atom is the Helium atom, its radius is 32 picometers. The largest atom is the cesium atom, which has a radius of 225 picometers. The prefix pico means ten to the minus twelfth power! That is, if we reduce 32 meters by a thousand billion times, we get the size of the radius of a helium atom.
At the same time, the scale of things is such that, in fact, the atom is 99% empty. The nucleus and electrons occupy an extremely small part of its volume. For clarity, consider this example. If you imagine an atom in the form of the Olympic stadium in Beijing (or maybe not in Beijing, just imagine a large stadium), then the nucleus of this atom will be a cherry located in the center of the field. The electron orbits would be somewhere at the level of the upper stands, and the cherry would weigh 30 million tons. Impressive, isn't it?
Where do atoms come from?
As you know, various atoms are now grouped in the periodic table. It contains 118 (and if with predicted but not yet discovered elements - 126) elements, not counting isotopes. But this was not always the case.
At the very beginning of the formation of the Universe, there were no atoms, and even more so, there were only elementary particles that interacted with each other under the influence of enormous temperatures. As a poet would say, it was a real apotheosis of particles. In the first three minutes of the existence of the Universe, due to a decrease in temperature and the coincidence of a whole bunch of factors, the process of primary nucleosynthesis began, when the first elements appeared from elementary particles: hydrogen, helium, lithium and deuterium (heavy hydrogen). It was from these elements that the first stars were formed, in the depths of which thermonuclear reactions took place, as a result of which hydrogen and helium “burned”, forming heavier elements. If the star was large enough, then it ended its life with a so-called “supernova” explosion, as a result of which atoms were thrown into the surrounding space. This is how the entire periodic table turned out.
So, we can say that all the atoms that we are made of were once part of ancient stars.
Why doesn't the nucleus of an atom decay?
There are four types in physics fundamental interactions between particles and the bodies they compose. These are strong, weak, electromagnetic and gravitational interactions.
Thanks to strong interaction, which manifests itself on the scale of atomic nuclei and is responsible for the attraction between nucleons, the atom is such a “tough nut to crack.”
Not so long ago, people realized that when the nuclei of atoms split, enormous energy was released. The fission of heavy atomic nuclei is a source of energy in nuclear reactors and nuclear weapons.
So, friends, having introduced you to the structure and basics of the structure of the atom, we can only remind you that we are ready to come to your aid at any time. It doesn’t matter whether you need to complete a diploma in nuclear physics, or the smallest test - situations are different, but there is a way out of any situation. Think about the scale of the Universe, order work from Zaochnik and remember - there is no reason to worry.
Since when chemical reactions the nuclei of reacting atoms remain unchanged (except for radioactive transformations), then chemical properties atoms depend on the structure of their electron shells. Theory electronic structure of the atom built on the basis of the apparatus of quantum mechanics. Thus, the structure of the energy levels of an atom can be obtained on the basis of quantum mechanical calculations of the probabilities of finding electrons in the space around the atomic nucleus ( rice. 4.5).
Rice. 4.5. Scheme of dividing energy levels into sublevels
The fundamentals of the theory of the electronic structure of an atom are reduced to the following provisions: the state of each electron in an atom is characterized by four quantum numbers: the principal quantum number n = 1, 2, 3,; orbital (azimuthal) l=0,1,2, n–1; magnetic m l = –l, –1,0,1, l; spin m s = -1/2, 1/2 .
According to Pauli principle, in the same atom there cannot be two electrons having the same set of four quantum numbers n, l, m l , m s; collections of electrons with the same principal quantum numbers n form electron layers, or energy levels of the atom, numbered from the nucleus and denoted as K, L, M, N, O, P, Q, and in the energy layer with a given value n can be no more than 2n 2 electrons. Collections of electrons with the same quantum numbers n And l, form sublevels, designated as they move away from the core as s, p, d, f.
The probabilistic determination of the position of the electron in space around the atomic nucleus corresponds to the Heisenberg uncertainty principle. According to quantum mechanical concepts, an electron in an atom does not have a specific trajectory of motion and can be located in any part of the space around the nucleus, and its various positions are considered as an electron cloud with a certain negative charge density. The space around the nucleus in which an electron is most likely to be found is called orbital. It contains about 90% of the electron cloud. Each sublevel 1s, 2s, 2p etc. corresponds to a certain number of orbitals of a certain shape. For example, 1s- And 2s- orbitals are spherical and 2p-orbitals ( 2p x , 2p y , 2p z-orbitals) are oriented in mutually perpendicular directions and have the shape of a dumbbell ( rice. 4.6).
Rice. 4.6. Shape and orientation of electron orbitals.
During chemical reactions, the atomic nucleus does not undergo changes, only electronic shells atoms, the structure of which explains many properties of chemical elements. Based on the theory of the electronic structure of the atom, the deep physical meaning of Mendeleev’s periodic law of chemical elements was established and the theory of chemical bonding was created.
The theoretical justification of the periodic system of chemical elements includes data on the structure of the atom, confirming the existence of a connection between the periodicity of changes in the properties of chemical elements and the periodic repetition of similar types of electronic configurations of their atoms.
In the light of the doctrine of the structure of the atom, Mendeleev’s division of all elements into seven periods becomes justified: the number of the period corresponds to the number of energy levels of atoms filled with electrons. In small periods, with an increase in the positive charge of atomic nuclei, the number of electrons at the external level increases (from 1 to 2 in the first period, and from 1 to 8 in the second and third periods), which explains the change in the properties of elements: at the beginning of the period (except for the first) there is alkali metal, then a gradual weakening of metallic properties and strengthening of non-metallic properties is observed. This pattern can be traced for elements of the second period in table 4.2.
Table 4.2.
IN long periods As the charge of nuclei increases, filling levels with electrons becomes more difficult, which explains the more complex change in the properties of elements compared to elements of small periods.
The identical nature of the properties of chemical elements in subgroups is explained by the similar structure of the external energy level, as shown in table 4.3, illustrating the sequence of filling energy levels with electrons for subgroups of alkali metals.
Table 4.3.
The group number usually indicates the number of electrons in an atom that can participate in the formation of chemical bonds. This is the physical meaning of the group number. In four places of the periodic table, the elements are not arranged in order of increasing atomic mass: Ar And K,Co And Ni,Te And I,Th And Pa. These deviations were considered shortcomings of the periodic table of chemical elements. The doctrine of the structure of the atom explained these deviations. Experimental determination of nuclear charges showed that the arrangement of these elements corresponds to an increase in the charges of their nuclei. In addition, the experimental determination of the charges of atomic nuclei made it possible to determine the number of elements between hydrogen and uranium, as well as the number of lanthanides. Now all places in the periodic table are filled in the interval from Z=1 to Z=114, however, the periodic system is not complete, the discovery of new transuranium elements is possible.
