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The distribution of electrons over various AOs is called electronic configuration of an atom. Electronic configuration with the lowest energy corresponds to basic state atom, the remaining configurations refer to excited states.
The electronic configuration of an atom is depicted in two ways - in the form of electronic formulas and electron diffraction diagrams. When writing electronic formulas, the principal and orbital quantum numbers are used. The sublevel is denoted by the principal quantum number (number) and the orbital quantum number (corresponding letter). The number of electrons in a sublevel characterizes the superscript. For example, for the ground state of the hydrogen atom, the electronic formula is: 1 s 1 .
The structure of electronic levels can be described more completely using electron diffraction diagrams, where the distribution over sublevels is represented in the form of quantum cells. In this case, the orbital is conventionally depicted as a square, near which the sublevel designation is affixed. The sub-levels at each level should be slightly offset in height, since their energy is somewhat different. Electrons are represented by arrows or ↓ depending on the sign of the spin quantum number. Electron diffraction diagram of the hydrogen atom:
The principle of constructing the electronic configurations of multielectron atoms is to add protons and electrons to the hydrogen atom. The distribution of electrons over energy levels and sublevels obeys the previously discussed rules: the principle least energy, the Pauli principle and Hund's rule.
Taking into account the structure of electronic configurations of atoms, all known elements, in accordance with the value of the orbital quantum number of the last filled sublevel, can be divided into four groups: s-elements, p-elements, d-elements, f-elements.
In a helium atom He (Z=2) the second electron occupies 1 s-orbital, its electronic formula: 1 s 2. Electronographic diagram:
Helium ends the first shortest period of the Periodic Table of the Elements. The electronic configuration of helium is denoted .
The second period opens lithium Li (Z=3), its electronic formula: 
Electronographic diagram:
The following are simplified electron diffraction diagrams of atoms of elements whose orbitals of the same energy level are located at the same height. Internal, fully filled sublevels are not shown.
Lithium is followed by beryllium Be (Z=4), in which an additional electron populates 2 s-orbital. Electronic formula Be: 2 s 2
In the ground state, the next boron electron B (z=5) occupies 2 R-orbital, V:1 s 2 2s 2 2p one ; its electron diffraction pattern:
The following five elements have electronic configurations:
C (Z=6): 2 s 2 2p 2N (Z=7): 2 s 2 2p 3
O (Z=8): 2 s 2 2p 4 F (Z=9): 2 s 2 2p 5
Ne (Z=10): 2 s 2 2p 6
The given electronic configurations are determined by Hund's rule.
The first and second energy levels of neon are completely filled. Let's designate its electronic configuration and we will use further for brevity of record of electronic formulas of atoms of elements.
Sodium Na (Z=11) and Mg (Z=12) open the third period. Outer electrons occupy 3 s-orbital:
Na (Z=11): 3 s 1
Mg (Z=12): 3 s 2
Then, starting with aluminum (Z=13), 3 R-sublevel. The third period ends with argon Ar (Z=18):
Al (Z=13): 3 s 2 3p 1
Ar (Z=18): 3 s 2 3p 6
The elements of the third period differ from the elements of the second in that they have free 3 d-orbitals that can participate in the formation of a chemical bond. This explains the valence states exhibited by the elements.
In the fourth period, in accordance with the rule ( n+l), in potassium K (Z=19) and calcium Ca (Z=20) electrons occupy 4 s- sublevel, not 3 d. Starting with scandium Sc (Z=21) and ending with zinc Zn (Z=30), 3 d- sublevel:
Electronic formulas d-elements can be represented in ionic form: sublevels are listed in ascending order of the main quantum number, and at a constant n– in order of increasing orbital quantum number. For example, for Zn such an entry would look like this: 
Both of these entries are equivalent, but the zinc formula given earlier correctly reflects the order in which the sublevels are filled.
Row 3 d-elements in chromium Cr (Z=24) there is a deviation from the rule ( n+l). According to this rule, the Cr configuration should look like this: 
Its real configuration is found to be 
Sometimes this effect is called the "failure" of the electron. Similar effects are explained by the increased stability by half ( p
3 ,
d
5 ,
f
7) and completely ( p
6 ,
d
10 ,
f
14) completed sublevels.
Deviations from the rule ( n+l) are also observed in other elements (Table 2). This is due to the fact that as the principal quantum number increases, the differences between the energies of the sublevels decrease.
Next comes filling 4 p-sublevel (Ga - Kr). The fourth period contains only 18 elements. Similarly, filling 5 s-, 4d- and 5 p- sublevels of 18 elements of the fifth period. Note that the energy 5 s- and 4 d-sublevels are very close, and an electron with 5 s- sub-level can easily go to 4 d-sublevel. On 5 s-sublevel Nb, Mo, Tc, Ru, Rh, Ag has only one electron. In basic condition 5 s- sublevel Pd is not filled. A “dip” of two electrons is observed.
table 2
Exceptions from ( n+l) – rules for the first 86 elements
|
Electronic configuration |
||
|
according to the rule ( n+l) |
actual |
|
|
4s 2 3d 4 4s 2 3d 9 5s 2 4d 3 5s 2 4d 4 5s 2 4d 5 5s 2 4d 6 5s 2 4d 7 5s 2 4d 8 5s 2 4d 9 6s 2 4f 1 5d 0 6s 2 4f 2 5d 0 6s 2 4f 8 5d 0 6s 2 4f 14 5d 7 6s 2 4f 14 5d 8 6s 2 4f 14 5d 9 |
4s 1 3d 5 4s 1 3d 10 5s 1 4d 4 5s 1 4d 5 5s 1 4d 6 5s 1 4d 7 5s 1 4d 8 5s 0 4d 10 5s 1 4d 10 6s 2 4f 0 5d 1 6s 2 4f 1 5d 1 6s 2 4f 7 5d 1 6s 0 4f 14 5d 9 6s 1 4f 14 5d 9 6s 1 4f 14 5d 10 |
|
In the sixth period after filling 6 s-sublevel of cesium Cs (Z=55) and barium Ba (Z=56) the next electron, according to the rule ( n+l), should take 4 f-sublevel. However, in lanthanum La (Z=57), an electron enters 5 d-sublevel. Half filled (4 f 7) 4f-sublevel has increased stability, therefore, gadolinium Gd (Z=64), following europium Eu (Z=63), by 4 f-sublevel retains the previous number of electrons (7), and the new electron arrives at 5 d-sublevel, breaking the rule ( n+l). In terbium Tb (Z=65), the next electron occupies 4 f-sublevel and there is an electron transition from 5 d- sublevel (configuration 4 f 9 6s 2). Filling 4 f-sublevel ends at ytterbium Yb (Z=70). The next electron of the lutetium atom Lu occupies 5 d-sublevel. Its electronic configuration differs from that of the lanthanum atom only by being completely filled with 4 f-sublevel.
Currently in Periodic system elements D.I. Mendeleev, under scandium Sc and yttrium Y, lutetium (rather than lanthanum) is sometimes located as the first d-element, and all 14 elements in front of it, including lanthanum, putting it in a special group lanthanides beyond the Periodic Table of the Elements.
The chemical properties of elements are determined mainly by the structure of the outer electronic levels. Change in the number of electrons on the third outside 4 f- sublevel has little effect on the chemical properties of elements. So all 4 f elements are similar in their properties. Then in the sixth period there is a filling of 5 d-sublevel (Hf - Hg) and 6 p-sublevel (Tl - Rn).
In the seventh period 7 s-sublevel is filled for francium Fr (Z=87) and radium Ra (Z=88). Actinium has a deviation from the rule ( n+l), and the next electron populates 6 d- sublevel, not 5 f. This is followed by a group of elements (Th - No) with a filling 5 f-sublevels that form a family actinides. Note that 6 d- and 5 f- sublevels have such close energies that the electronic configuration of actinide atoms often does not obey the rule ( n+l). But in this case fine configuration value 5 f T 5d m not so important, since it has a rather small effect on Chemical properties element.
Lawrencium Lr (Z=103) has a new electron at 6 d-sublevel. This element is sometimes placed in the Periodic Table under lutetium. The seventh period is not completed. Elements 104 – 109 are unstable and their properties are little known. Thus, as the charge of the nucleus increases, similar electronic structures of the outer levels are periodically repeated. In this regard, one should also expect periodic changes in various properties of elements.
Periodic change in the properties of atoms of chemical elements
The chemical properties of the atoms of elements are manifested during their interaction. The types of configurations of the external energy levels of atoms determine the main features of their chemical behavior.
The characteristics of the atom of each element that determine its behavior in chemical reactions are the ionization energy, electron affinity, electronegativity.
Ionization energy is the energy required to detach and remove an electron from an atom. The lower the ionization energy, the higher the reducing power of the atom. Therefore, the ionization energy is a measure of the reducing ability of an atom.
The ionization energy required to detach the first electron is called the first ionization energy I 1 . The energy required to detach the second electron is called the second ionization energy I 2, etc. In this case, the following inequality takes place
I 1< I 2 < I 3 .
The detachment and removal of an electron from a neutral atom occurs more easily than from a charged ion.
The maximum value of the ionization energy corresponds to the noble gases. Alkali metals have the minimum value of ionization energy.
Within one period, the ionization energy varies nonmonotonically. Initially, it decreases when moving from s-elements to the first p-elements. Then, in subsequent p-elements, it increases.
Within one group, with an increase in the ordinal number of the element, the ionization energy decreases, which is due to an increase in the distance between the outer level and the nucleus.
Electron affinity is the energy (denoted by E) that is released when an electron is attached to an atom. When an atom accepts an electron, it becomes a negatively charged ion. The electron affinity in a period increases, while in a group, as a rule, it decreases.
Halogens have the highest electron affinity. By attaching the missing electron to complete the shell, they acquire the completed configuration of a noble gas atom.
Electronegativity is the sum of ionization energy and electron affinity
Electronegativity increases with a period and decreases with a subgroup.
Atoms and ions do not have strictly defined boundaries due to the wave nature of the electron. Therefore, the radii of atoms and ions are determined conditionally.
The greatest increase in the radius of atoms is observed in elements of small periods, in which only the outer energy level is filled, which is typical for s- and p-elements. For d- and f-elements, a smoother increase in the radius with increasing nuclear charge is observed.
Within a subgroup, the atomic radius increases as the number of energy levels increases.
>> Chemistry: Electronic configurations of atoms chemical elements
The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons that have opposite (antiparallel) spins (translated from English as “spindle”), that is, they have properties that can be conditionally represented itself as the rotation of an electron around its imaginary axis: clockwise or counterclockwise. This principle is called the Pauli principle.
If there is one electron in the orbital, then it is called unpaired, if there are two, then these are paired electrons, that is, electrons with opposite spins.
Figure 5 shows a diagram of the division of energy levels into sublevels.
The s-orbital, as you already know, is spherical. The electron of the hydrogen atom (s = 1) is located on this orbital and is unpaired. Therefore, its electronic formula or electronic configuration will be written as follows: 1s 1. In electronic formulas, the energy level number is indicated by the number in front of the letter (1 ...), the sublevel (orbital type) is indicated by the Latin letter, and the number that is written to the upper right of the letter (as an exponent) shows the number of electrons in the sublevel.
For a helium atom, He, having 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, since they are at a greater distance from the nucleus than 1s-orbital electrons (n = 2).
In general, for every value of n, there is one s-orbital, but with a corresponding amount of electron energy in it and, therefore, with a corresponding diameter, growing as the value of n increases.
p-Orbital 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, and z axes.
For elements of the second period (n = 2), first one β-orbital is filled, and then three p-orbitals. Electronic formula 1l: 1s 2 2s 1. The electron is weaker bound to the nucleus of the atom, so the lithium atom can easily give it away (as you probably remember, this process is called oxidation), turning into a Li + ion.
In the beryllium atom Be 0, the fourth electron is also located in the 2s orbital: 1s 2 2s 2 . The two outer electrons of the beryllium atom are easily detached - Be 0 is oxidized to the Be 2+ cation.
At the boron atom, the fifth electron occupies a 2p orbital: 1s 2 2s 2 2p 1. Further, the atoms C, N, O, E are filled with 2p orbitals, which ends with the noble gas neon: 1s 2 2s 2 2p 6.
For the elements of the third period, the Sv- and Sp-orbitals are filled, respectively. Five d-orbitals of the third level remain free:
11 Na 1s 2 2s 2 Sv1; 17C11v22822r63r5; 18Ar P^Yor^3p6.
Sometimes in diagrams depicting the distribution of electrons in atoms, only the number of electrons at each energy level is indicated, that is, they write down the abbreviated electronic formulas of atoms of chemical elements, in contrast to the full electronic formulas given above.
For elements of large periods (fourth and fifth), the first two electrons occupy the 4th and 5th orbitals, respectively: 19 K 2, 8, 8, 1; 38 Sr 2, 8, 18, 8, 2. Starting with the third element of each long 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 Tr 2, 8, 14, 2; 40 Zr 2, 8, 18, 10, 2; 43 Tr 2, 8, 18, 13, 2. As a rule, when the previous d-sublevel is filled, the outer (4p- and 5p, respectively) p-sublevel will begin to fill.
For elements of large periods - the sixth and incomplete seventh - electronic levels and sublevels are filled with electrons, as a rule, as follows: the first two electrons will go to the outer β-sublevel: 56 Ba 2, 8, 18, 18, 8, 2; 87Gr 2, 8, 18, 32, 18, 8, 1; the next one electron (for Na and Ac) to the previous (p-sublevel: 57 La 2, 8, 18, 18, 9, 2 and 89 Ac 2, 8, 18, 32, 18, 9, 2.
Then the next 14 electrons will go to the third energy level from the outside in the 4f and 5f orbitals, respectively, for lanthanides and actinides.
Then the second outside energy level (d-sublevel) will begin to build up again: for elements of secondary subgroups: 73 Ta 2, 8.18, 32.11, 2; 104 Rf 2, 8.18, 32, 32.10, 2, - and, finally, only after complete filling with ten electrons of the current level will the outer p-sublevel 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: the Pauli principle, according to which there can be no more than two electrons in a cell (orbitals, but with antiparallel spins), and F. Hund's rule, according to which electrons occupy free cells (orbitals), are located in they are first one at a time and at the same time have the same spin value, and only then they pair, but the spins in this case, according to the Pauli principle, will already be oppositely directed.
In conclusion, let us once again consider the mapping of the electronic configurations of atoms of elements over the periods of the D. I. Mendeleev system. Schemes of the electronic structure of atoms show the distribution of electrons over electronic layers (energy levels). 
In a helium atom, the first electron layer is completed - it has 2 electrons.
Hydrogen and helium are s-elements; these atoms have an s-orbital 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 e- 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 (Table 2).
In the neon atom, the second electron layer is completed - it has 8 electrons.
Table 2 The structure of the electron shells of atoms of elements of the second period
The end of the table. 2 
Li, Be - in-elements.
B, C, N, O, F, Ne - p-elements, these atoms are filled with electrons p-orbitals.
Elements of the third period
For atoms of elements of the third period, the first and second electron layers are completed; therefore, the third electron layer is filled, in which electrons can occupy the 3s, 3p, and 3d sublevels (Table 3).
Table 3 The structure of the electron shells of atoms of elements of the third period
A 3s-electron orbital is completed at the magnesium atom. Na and Mg-s elements. 
There are 8 electrons in the outer layer (the third electron layer) in the argon atom. As an outer layer, it is complete, but in total, in the third electron layer, as you already know, there can be 18 electrons, which means that the elements of the third period have unfilled 3d orbitals.
All elements from Al to Ag are p-elements. s- and p-elements form the main subgroups in the Periodic system.
A fourth electron layer appears at the potassium and calcium atoms, and the 4s sublevel is filled (Table 4), since it has a lower energy than the 3d sublevel. To simplify the graphical electronic formulas of the atoms of the elements of the fourth period: 1) we denote the conditionally graphical electronic formula of argon as follows:
Ar;
2) we will not depict the sublevels that are not filled for these atoms.
Table 4 The structure of the electron shells of atoms of the elements of the fourth period
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 the secondary subgroups, they have a pre-external electron layer filled, they are referred to as transition elements.
Pay attention to the structure of the electron shells of chromium and copper atoms. In them, a "failure" of one electron from the 4n- to the 3d sublevel occurs, which is explained by the greater energy stability of the resulting electronic configurations 3d 5 and 3d 10: 
In the zinc atom, the third electron layer is completed - all the 3s, 3p and 3d sublevels are filled in it, in total there are 18 electrons on them.
In the elements following zinc, the fourth electron layer continues to be filled, the 4p sublevel: Elements from Ga to Kr are p-elements. 
The outer layer (fourth) of the krypton atom is complete and has 8 electrons. But just in the fourth electron layer, as you know, there can be 32 electrons; the 4d and 4f sublevels of the krypton atom still remain unfilled.
The elements of the fifth period are filling the sublevels in the following order: 5s-> 4d -> 5p. And there are also exceptions associated with the "failure" of electrons, in 41 Nb, 42 MO, etc.
In the sixth and seventh periods, elements appear, that is, elements in which the 4f and 5f sublevels of the third outside electronic layer are being filled, respectively.
The 4f elements are called lanthanides.
5f-elements are called actinides.
The order of filling of electronic sublevels in the atoms of elements of the sixth period: 55 Сs and 56 Ва - 6s-elements;
57 La... 6s 2 5d 1 - 5d element; 58 Ce - 71 Lu - 4f elements; 72 Hf - 80 Hg - 5d elements; 81 Tl- 86 Rn - 6p-elements. But even here there are elements in which the order of filling is “violated” electron orbitals, which, for example, is associated with greater energy stability of half and fully filled f sublevels, that is, nf 7 and nf 14 .
Depending on which sublevel of the atom is filled with electrons last, all elements, as you already understood, are divided into four electronic families or blocks (Fig. 7).
1) s-Elements; the β-sublevel of the outer level of the atom is filled with electrons; s-elements include hydrogen, helium and elements of the main subgroups of groups I and II;
2) p-elements; the p-sublevel of the outer level of the atom is filled with electrons; p elements include elements of the main subgroups of III-VIII groups;
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, that is, elements of intercalary decades of large periods located between s- and p-elements. They are also called transition elements;
4) f-elements, the f-sublevel of the third outside level of the atom is filled with electrons; these include lanthanides and actinides.
1. What would happen if the Pauli principle was not respected?
2. What would happen if Hund's rule was not respected?
3. Make diagrams of the electronic structure, electronic formulas and graphic electronic formulas of atoms of the following chemical elements: Ca, Fe, Zr, Sn, Nb, Hf, Ra.
4. Write the electronic formula for element #110 using the symbol for the corresponding noble gas.
Lesson content lesson summary support frame lesson presentation accelerative methods interactive technologies Practice tasks and exercises self-examination workshops, trainings, cases, quests homework discussion questions rhetorical questions from students Illustrations audio, video clips and multimedia photographs, pictures graphics, tables, schemes humor, anecdotes, jokes, comics parables, sayings, crossword puzzles, quotes Add-ons abstracts articles chips for inquisitive cheat sheets textbooks basic and additional glossary of terms other Improving textbooks and lessonscorrecting errors in the textbook updating a fragment in the textbook elements of innovation in the lesson replacing obsolete knowledge with new ones Only for teachers perfect lessons calendar plan for the year guidelines discussion programs Integrated LessonsThe structure of the electron shells of atoms of the elements of the first four periods: $s-$, $p-$ and $d-$elements. The electronic configuration of the atom. Ground and excited states of atoms
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, based on experiments, came to the conclusion that electricity is transferred tiny 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), which differ somewhat 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; 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:
- $α$-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;
- $β$-rays are a stream of electrons;
- $γ$-rays - electromagnetic waves with a negligible mass, not carrying 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 serial number 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:
- we denote conditionally the graphic electronic formula of argon as follows: $Ar$;
- 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:
- $s$ -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 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;
- $f$ -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.
Electronic configuration- the formula for the arrangement of electrons in various electron shells of an atom of a chemical element or molecule.
The electronic configuration is usually written for atoms in their ground state. To determine the electronic configuration of an element, the following rules exist:
- Filling principle. According to the filling principle, electrons in the ground state of an atom fill orbitals in a sequence of increasing orbital energy levels. The lowest energy orbitals are always filled first.
- Pauli exclusion principle. According to this principle, no more than two electrons can be in any orbital, and then only if they have opposite spins (unequal spin numbers).
- Hund's rule. According to this rule, the filling of the orbitals of one subshell begins with single electrons with parallel (same in sign) spins, and only after single electrons have occupied all the orbitals, the final filling of the orbitals with pairs of electrons with opposite spins can occur.
From the point of view of quantum mechanics, the electronic configuration is a complete list of one-electron wave functions , from which, with a sufficient degree of accuracy, it is possible to compose the complete wave function of an atom (in the approximation of a self-consistent field).
Generally speaking, the atom, as a composite system, can only be fully described by the full wave function. However, such a description is practically impossible for atoms more complex than the hydrogen atom, the simplest of all atoms of chemical elements. A convenient approximate description is the self-consistent field method. This method introduces the concept of the wave function of each electron. The wave function of the entire system is written as a properly symmetrized product of one-electron wave functions. When calculating the wave function of each electron, the field of all other electrons is taken into account as an external potential, which in turn depends on the wave functions of these other electrons.
As a result of applying the self-consistent field method, we obtain a complex system nonlinear integro-differential equations , which is still difficult to solve. However, the self-consistent field equations have the rotational symmetry of the original problem (that is, they are spherically symmetric). This makes it possible to completely classify the one-electron wave functions that make up the complete wave function of an atom.
To begin with, as in any centrally symmetric potential, the wave function in a self-consistent field can be characterized by the quantum number of the total angular momentum l (\displaystyle l) and the quantum number of the projection of the angular momentum on some axis m (\displaystyle m). Wave functions with different values m (\displaystyle m) correspond to the same energy level, i.e., they are degenerate. Also, one energy level corresponds to states with different projections of the electron spin on any axis. Total for a given energy level 2 (2 l + 1) (\displaystyle 2(2l+1)) wave functions. Further, for a given value of the angular momentum, the energy levels can be renumbered. By analogy with the hydrogen atom, it is customary to number the energy levels for a given l (\displaystyle l) beginning with n = l + 1 (\displaystyle n=l+1). Full list quantum numbers of one-electron wave functions, from which the wave function of an atom can be composed, and is called the electronic configuration. Since everything is degenerate in quantum number m (\displaystyle m) and on the spin, it is enough to indicate the total number of electrons in the state with the data n (\displaystyle n), l (\displaystyle l).
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For historical reasons, in the electronic configuration formula, the quantum number l (\displaystyle l) written in Latin letters. The state with is denoted by the letter s (\displaystyle s), p (\displaystyle p): l = 1 (\displaystyle l=1), d (\displaystyle d): l = 2 (\displaystyle l=2), f (\displaystyle f): l = 3 (\displaystyle l=3), g (\displaystyle g): l = 4 (\displaystyle l=4) and so on alphabetically. To the left of the number l (\displaystyle l) write a number n (\displaystyle n), and on top of the number l (\displaystyle l) is the number of electrons in the data state n (\displaystyle n) and l (\displaystyle l). for instance 2 s 2 (\displaystyle 2s^(2)) corresponds to two electrons in a state with n = 2 (\displaystyle n=2), l = 0 (\displaystyle l=0). Due to practical convenience (see the Klechkovsky rule), in the full formula for the electronic configuration, the terms are written in ascending order of the quantum number n (\displaystyle n), and then the quantum number l (\displaystyle l), For example 1 s 2 2 s 2 2 p 6 3 s 2 3 p 3 (\displaystyle 1s^(2)2s^(2)2p^(6)3s^(2)3p^(3)). Since such a notation is somewhat redundant, sometimes the formula is reduced to 1 s 2 2 s 2 p 6 3 s 2 p 3 (\displaystyle 1s^(2)2s^(2)p^(6)3s^(2)p^(3)), i.e. omit the number n (\displaystyle n) where it can be guessed from the term ordering rule.
Periodic law and the structure of the atom
All those involved in the structure of the atom in any of their studies proceed from the tools that are provided to them by the periodic law, discovered by the chemist D. I. Mendeleev; only in their understanding of this law, physicists and mathematicians use their “language” to interpret the dependencies shown by him (although J. W. Gibbs is known for a rather ironic aphorism on this subject), but, at the same time, isolated from chemists studying matter , with all the perfection, advantages and universality of their apparatus, neither physicists nor mathematicians, of course, can build their own research.
The interaction of representatives of these disciplines is also observed in further development themes. The discovery of secondary periodicity by E. V. Biron (1915) gave another aspect in understanding the issues related to the regularities of the structure of electron shells. S. A. Shukarev, student of E. V. Biron and
Lecture 2. Electronic configuration of the element
At the end of the last lecture, based on the Klechkovsky rules, we constructed the order of filling the energy sublevels with electrons
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 5d1 4f14 5d9 6p6 7s2 6d1 5f14 6d9 7p6 …
The distribution of electrons of an atom over energy sublevels is called electronic configuration. First of all, when looking at the filling series, a certain periodicity-regularity is striking.
The filling of energy orbitals with electrons in the ground state of an atom obeys the principle of least energy: first, more favorable low-lying orbitals are filled, and then sequentially higher-lying orbitals according to the order of filling.
Let's analyze the filling sequence.
If exactly 1 electron is present in an atom, it falls into the lowest-lying 1s-AO (AO is an atomic orbital). Therefore, the emerging electronic configuration can be represented by the notation 1s1 or graphically (See below - an arrow in a box).
It is easy to understand that if there is more than one electron in an atom, they sequentially occupy first 1s, and then 2s, and, finally, go to the 2p sublevel. However, already for six electrons (a carbon atom in the ground state), two possibilities arise: filling the 2p sublevel with two electrons with the same spin or with the opposite.
Let's use a simple analogy: suppose that atomic orbitals are a kind of "rooms" for "residents", in the role of which electrons act. From practice it is well known that tenants prefer to occupy each private room and not crowded into one.
A similar behavior is also characteristic of electrons, which is reflected in Hund's rule:
Hund's rule: the steady state of an atom corresponds to such a distribution of electrons within the energy sublevel at which the total spin is maximum.
The state of an atom with a minimum energy is called the ground state, and all the rest are called excited states of the atom.
Lecture 2. Electronic configuration
Atoms of elements of I and II periods
1 electron
2 electrons
3 electrons
4 electrons
5 electrons
6 electrons
7 electrons
8 electrons
9 electrons
10Ne
10 electrons
Element of the whole e-
electronic configuration
electron distribution
Then, based on Hund's rule, for nitrogen the ground state assumes the presence of three unpaired p-electrons (electron configuration …2p3 ). In the atoms of oxygen, fluorine, and neon, sequential pairing of electrons occurs and the 2p sublevel is filled.
Note that the third period of the Periodic Table begins with the sodium atom,
whose configuration (11 Na ... 3s1 ) is very similar to that of lithium (3 Li ... 2s1 )
except that the principal quantum number n is three instead of two.
The filling of energy sublevels in atoms of period III elements with electrons is exactly the same as that observed for elements of period II: the magnesium atom completes the filling of the 3s sublevel, then from aluminum to argon, electrons are sequentially placed on the 3p sublevel according to Hund's rule: first, individual electrons are placed on the AO ( Al, Si, P), then their pairing occurs.
Atoms of elements of the III period
11Na
12Mg
13Al
14Si
17Cl
18Ar
abbreviated
e- distribution
Lecture 2. Electronic configuration
The fourth period of the Periodic Table begins with the filling of the 4s-sublevel in potassium and calcium atoms with electrons. As follows from the order of filling, then comes the turn of 3d orbitals.
Thus, we can conclude that the filling of d-AO with electrons is “late” by 1 period: in the IV period, 3 (!) d-sublevel is filled).
So, from Sc to Zn, the 3d sublevel (10 electrons) is filled with electrons, then from Ga to Kr, the 4p sublevel is filled.
Atoms of elements of period IV
20Ca
21sc
1s2 2s2 2p6 3s2 3p6 4s2 3d1
4s2 3d1
1s2 2s2 2p6 3s2 3p6 4s2 3d2
22ti
4s2 3d2
30Zn
1s2 2s2 2p6 3s2 3p6 4s2 3d10
4s2 3d10
31Ga
1s 2s 2p 3s 3p 4s 3d
36Kr
1s 2s 2p 3s 3p 4s 3d
abbreviated
e- distribution
The filling of energy sublevels with electrons in the atoms of elements of period V is exactly the same as that observed for elements of period IV
(disassemble by yourself)
In the sixth period, the 6s sublevel is first filled with electrons (55 Cs atoms and
56 Ba), and then one electron is located on the 5d orbital of lanthanum (57 La 6s2 5d1 ).
For the next 14 elements (from 58 to 71), the 4f sublevel is filled, i.e. the filling of f-orbitals is “late” by 2 periods, while the electron at the 5d sublevel is preserved. For example, one should write down the electronic configuration of cerium
58 Ce 6s2 5d 1 4 f 1
Starting from element 72 (72 Hf) and up to element 80 (80 Hg), the 5d sublevel is “filled in”.
Therefore, the electronic configurations of hafnium and mercury are
72 Hf 6s2 5d 1 4 f 14 5d 1 or 72 Hf 6s2 4 f 14 5d 2 80 Hg 6s2 5d 1 4 f 14 5d 9 or 80 Hg 6s2 4 f 14 5d 10
Lecture 2. Electronic configuration
Similarly, the energy sublevels in the atoms of the elements of the VII period are filled with electrons.
Determination of quantum numbers from electronic configuration
What are quantum numbers, how did they appear and why are they needed - see Lecture 1.
Given: Electronic configuration entry "3p 4"
The principal quantum number n is the first digit in the entry, i.e. "3". n = 3 "3 p4", the main quantum number;
Side (orbital, azimuthal) quantum number l is encoded letter designation sublevel. The letter p corresponds to the number l = 1.
cloud shape
l \u003d 1 "3p 4",
"dumbbell"
Distribution of electrons within a sublevel according to the Pauli principle and Hund's rule
m Є [-1; +1] - the orbitals are the same (degenerate) in energy n = 3, l = 1, m Є [-1; +1] (m = -1); s = + ½
n = 3, l = 1, m Є [-1; +1] (m = 0); s = + ½n = 3, l = 1, m Є [-1; +1] (m = +1); s = + ½ n = 3, l = 1, m Є [-1; +1] (m = -1); s = -½
Valence level and valence electrons
Valence level called a set of energy sublevels that are involved in the formation chemical bonds with other atoms.
Valence electrons are those located at the valence level.
Elements of PSCE are divided into 4 groups
s-elements. Valence electrons ns x . The two s-elements are at the beginning of each period.
p-elements. Valence electrons ns 2 np x . Six p-elements are located at the end of each period (except the first and seventh).
Lecture 2. Electronic configuration
d-elements. Valence electrons ns 2 (n-1)d x. Ten d-elements form secondary subgroups, starting from period IV and are between s- and p-elements.
f-elements. Valence electrons ns 2 (n-1)d 1 (n-2)f x. The fourteen f-elements form the series of lanthanides (4f) and actinides (5f), which are located below the table.
Electronic analogues are particles that are characterized by similar electronic configurations, i.e. distribution of electrons over sublevels.
for instance
H 1s1 Li … 2s1 Na … 3s1 K … 4s1
Electronic analogues have similar electronic configurations, so their chemical properties are similar - and they are located in the Periodic system of elements in the same subgroup.
Electronic "failure" (or electronic "overshoot")
Quantum mechanics predicts that the state of the particle has the lowest energy when all levels are filled with electrons either completely or half.
So for chromium subgroup elements(Cr, Mo, W, Sg) and copper subgroup elements(Cu, Ag, Au) there is a displacement of 1 electron from s - to the d- sublevel.
24 Cr 4s2 3d4 24 Cr 4s1 3d5 29 Cu 4s2 3d9 29 Cu 4s1 3d10
This phenomenon is called electronic "failure", it should be remembered.
A similar phenomenon is also characteristic of f-elements, but their chemistry is beyond the scope of our course.
Please note: for p-elements, the electronic dip is NOT observed!
Summing up, it should be concluded that the number of electrons in an atom is determined by the composition of its nucleus, and their distribution (electronic configuration) is determined by sets
Lecture 2. Electronic configuration
quantum numbers. In turn, the electronic configuration determines the chemical properties of the element.
Therefore, it is obvious that Properties of simple substances, as well as properties of compounds
elements are in a periodic dependence on the magnitude of the charge of the nucleus
atom (serial number).
Periodic Law
Basic properties of atoms of elements
1. The radius of an atom is the distance from the center of the nucleus to the outer energy level. V
period, as the charge of the nucleus increases, the radius of the atom decreases; in Group,
on the contrary, as the number of energy levels increases, the radius of the atom increases.
Consequently, in the series O2- , F- , Ne, Na+ , Mg2+ - the particle radius decreases, although their configuration is the same 1s2 2s2 2p6 .
For non-metals, they talk about the covalent radius, for metals, about the metallic radius, for ions, about the ionic radius.
2. The ionization potential is the energy that must be spent on separation from atom 1
electron. According to the principle of least energy, the last electron in terms of filling (for s and p elements) and the electron of the outer energy level (for d and f elements) come off first
In the period, as the charge of the nucleus increases, the ionization potential grows - at the beginning of the period there is an alkali metal with a low ionization potential, at the end of the period - an inert gas. In a group, ionization potentials weaken.
Ionization energy, eV
3. Electron affinity - the energy released when an electron is attached to an atom, i.e. in the formation of an anion.
4. Electronegativity (EO) is the ability of atoms to attract electron density towards themselves. Unlike the ionization potential, which is followed by a specific measurable physical quantity, EO is a certain quantity that can beonly calculated, it cannot be measured. In other words, EO was invented by people in order to use it to explain certain phenomena.
For our educational purposes, it is required to remember the qualitative order of change
electronegativity: F > O > N > Cl > ... > H > ... > metals.
EO - the ability of an atom to shift its electron density towards itself, - obviously,
increases in the period (since the charge of the nucleus increases - the force of attraction of the electron and the radius of the atom decreases) and, on the contrary, weakens in the group.
It is easy to understand that since the period begins with an electropositive metal,
and ends with a typical non-metal of group VII (inert gases are not taken into account), then the degree of change in EC in the period is greater than in the group.
Lecture 2. Electronic configuration
5. The oxidation state is the conditional charge of an atom in a chemical compound,
calculated in the approximation that all bonds are formed by ions. The minimum oxidation state is determined by how many electrons an atom can accept per
represent the sequence in which atoms are connected to each other. Consider separately each pair of atoms and denote by an arrow the displacement of electrons to that atom from the pair, whose EC is greater than (b). Consequently, the electrons shifted - and charges were formed - positive and negative:
at the end of each arrow is a charge (-1), corresponding to the addition of 1 electron;
on the base of the arrow is the charge (+1) corresponding to the removal of 1 electron.
The resulting charges are the oxidation state of a particular atom.
H+1
H+1
That's all for today, thank you for your attention.
Literature
1. S.G. Baram, M.A. Ilyin. Chemistry at the Summer School. Proc. allowance / Novosib. state
un-t, Novosibirsk, 2012. 48 p.
2. A.V. Manuilov, V.I. Rodionov. Fundamentals of chemistry for children and adults. – M.:
CJSC Publishing House Tsentrpoligraf, 2014. - 416 p. - see p. 29-85. http://www.hemi.nsu.ru/
