s-Elements elements are called, in the atoms of which the last electron enters the s-sublevel. The p-elements,d-elements andf-elements.
The beginning of each period corresponds to the opening of a new electronic layer. The period number is equal to the number of the opened electron layer. Each period, except the first, ends with the filling of the p-sublevel of the layer opened at the beginning of this period. The first period contains only s-elements (two). In the fourth and fifth periods between the s- (two) and p-elements (six) are d-elements (ten). In the sixth and seventh, after a pair of s-elements, there is (in violation of Klechkovsky's rules) one d-element, then fourteen f-elements (they are placed in separate lines at the bottom of the table - lanthanides and actinides), then nine d-elements and, as always , the periods end with six p-elements.
Vertically, the table is divided into 8 groups, each group - into the main and secondary subgroups. In the main subgroups are s- and p-elements, in the secondary - d-elements. The main subgroup is easy to determine - it contains elements of 1-3 periods. Strictly below them are the remaining elements of the main subgroup. The elements of the secondary subgroup are located to the side (left or right).
Valence of atoms
In the classical view, valence is determined by the number of unpaired electrons in the ground or excited state of atoms. Basic state- electronic state of an atom in which its energy is minimal. excited state- the electronic state of an atom, corresponding to the transition of one or more electrons from an orbital with a lower energy to a free orbital with a higher energy. For s- and p-elements, the transition of electrons is possible only within the outer electronic layer. For d-elements, transitions are possible within the d-sublevel of the pre-outer layer and the s- and p-sublevels of the outer layer. For f-elements, transitions are possible within (n-2)f-, (n-1)d-, ns- and np-sublevels, where n is the number of the outer electronic layer. Valence electrons called electrons that determine the valency of an atom in its ground or excited state. Valence electron layer- the layer on which the valence electrons are located.
Describe using quantum numbers the electrons of the outer layer of the sulfur atom and the valence electrons of iron (ground state). Indicate the possible valencies and oxidation states of the atoms of these elements.
one). Sulfur atom.
Sulfur has serial number 16. It is in the third period, the sixth group, the main subgroup. Therefore, this is a p-element, the outer electronic layer is the third, and it is valence. It has six electrons. The electronic structure of the valence layer has the form
For all electrons n=3, since they are located on the third layer. Let's consider them in order:
n=3, L=0 (the electron is located in the s-orbital), m l =0 (for the s-orbital, only such a value of the magnetic quantum number is possible), m s =+1/2 (rotation around its own axis occurs clockwise) ;
n=3, L=0, m l \u003d 0 (these three quantum numbers are the same as those of the first electron, since both electrons are in the same orbital), m s \u003d -1/2 (only here the difference appears, required by the Pauli principle);
n=3, L=1 (this is a p-electron), m l \u003d +1 (out of three possible values m l \u003d 1, 0 for the first p-orbital, we choose the maximum, this is the p x-orbital), m s \u003d +1 / 2;
n=3, L=1, m l = +1, m s =-1/2;
n=3, L=1, m l \u003d 0 (this is r y-orbital), m s \u003d +1/2;
n=3, L=1, m l \u003d -1 (this is the p z-orbital), m s \u003d +1/2.
Consider the valencies and oxidation states of sulfur. On the valence layer in the ground state of the atom, there are two electron pairs, two unpaired electrons and five free orbitals. Therefore, the valency of sulfur in this state is II. Sulfur is a non-metal. Before the completion of the layer, it lacks two electrons, therefore, in compounds with atoms of less electronegative elements, for example, with metals, it can exhibit a minimum oxidation state of -2. Depairing of electron pairs is possible, because there are free orbitals on this layer. Therefore, in the first excited state (S*)
In compounds with atoms of more electronegative elements, such as oxygen, all six valence electrons can be displaced from sulfur atoms, so its maximum oxidation state is +6.
2). Iron.
The serial number of iron is 26. It is located in the fourth period, in the eighth group, a side subgroup. This is a d-element, the sixth in a series of d-elements of the fourth period. The valence electrons of iron (eight) are located on the 3d-sublevel (six, in accordance with the position in the row of d-elements) and on the 4s-sublevel (two):
Let's consider them in order:
n=3, L=2, m l = +2, m s = +1/2;
n=3, L=2, m l = +2, m s = -1/2;
n=3, L=2, m l = +1, m s = +1/2;
n=3, L=2, m l = 0, m s = +1/2;
n=3, L=2, m l = -1, m s = +1/2;
n=3, L=2, m l = -2, m s = +1/2;
n=4, L=0, m l = 0, m s = +1/2;
n=4, L=0, m l = 0, m s = -1/2.
Valence
There are no unpaired electrons on the outer layer, therefore, the minimum valency of iron (II) appears in the excited state of the atom:
After the electrons of the outer layer are used, 4 unpaired electrons of the 3d sublevel can be involved in the formation of chemical bonds. Therefore, the maximum valency of iron is VI.
Oxidation state
Iron is a metal, therefore, it is characterized by positive oxidation states from +2 (4s-sublevel electrons are involved) to +6 (4s- and all unpaired 3d-electrons are involved).
The location of electrons on energy shells or levels is recorded using electronic formulas of chemical elements. Electronic formulas or configurations help to represent the structure of an element's atom.
The structure of the atom
The atoms of all elements consist of a positively charged nucleus and negatively charged electrons that are located around the nucleus.
The electrons are at different energy levels. The farther an electron is from the nucleus, the more energy it has. The size of the energy level is determined by the size of the atomic orbit or orbital cloud. This is the space in which the electron moves.
Rice. 1. The general structure of the atom.
Orbitals can have different geometric configurations:
- s-orbitals- spherical;
- p-, d and f-orbitals- dumbbell-shaped, lying in different planes.
At the first energy level of any atom, there is always an s-orbital with two electrons (an exception is hydrogen). Starting from the second level, the s- and p-orbitals are at the same level.
Rice. 2. s-, p-, d and f-orbitals.
Orbitals exist regardless of the location of electrons on them and can be filled or vacant.
Formula entry
Electronic configurations of atoms of chemical elements are written according to the following principles:
- each energy level corresponds to a serial number, denoted by an Arabic numeral;
- the number is followed by a letter denoting the orbital;
- a superscript is written above the letter, corresponding to the number of electrons in the orbital.
Recording examples:
You need to enable JavaScript to run this app. Electronic configuration of an atom is a formula showing the arrangement of electrons in an atom by levels and sublevels. After studying the article, you will find out where and how electrons are located, get acquainted with quantum numbers and be able to build the electronic configuration of an atom by its number, at the end of the article there is a table of elements.
Why study the electronic configuration of elements?
Atoms are like a constructor: there are a certain number of parts, they differ from each other, but two parts of the same type are exactly the same. But this constructor is much more interesting than the plastic one, and here's why. The configuration changes depending on who is nearby. For example, oxygen next to hydrogen maybe turn into water, next to sodium into gas, and being next to iron completely turns it into rust. To answer the question why this happens and to predict the behavior of an atom next to another, it is necessary to study the electronic configuration, which will be discussed below.
How many electrons are in an atom?
An atom consists of a nucleus and electrons revolving around it, the nucleus consists of protons and neutrons. In the neutral state, each atom has the same number of electrons as the number of protons in its nucleus. The number of protons was indicated by the element's serial number, for example, sulfur has 16 protons - the 16th element of the periodic system. Gold has 79 protons - the 79th element of the periodic table. Accordingly, there are 16 electrons in sulfur in the neutral state, and 79 electrons in gold.
Where to look for an electron?
Observing the behavior of an electron, certain patterns were derived, they are described by quantum numbers, there are four of them in total:
- Principal quantum number
- Orbital quantum number
- Magnetic quantum number
- Spin quantum number
Orbital
Further, instead of the word orbit, we will use the term "orbital", the orbital is the wave function of the electron, roughly - this is the area in which the electron spends 90% of the time.
N - level
L - shell
M l - orbital number
M s - the first or second electron in the orbital
Orbital quantum number l
As a result of the study of the electron cloud, it was found that depending on the level of energy, the cloud takes four main forms: a ball, dumbbells and the other two, more complex. In ascending order of energy, these forms are called s-, p-, d- and f-shells. Each of these shells can have 1 (on s), 3 (on p), 5 (on d) and 7 (on f) orbitals. The orbital quantum number is the shell on which the orbitals are located. The orbital quantum number for s, p, d and f orbitals, respectively, takes the values 0,1,2 or 3.
On the s-shell one orbital (L=0) - two electrons
There are three orbitals on the p-shell (L=1) - six electrons
There are five orbitals on the d-shell (L=2) - ten electrons
There are seven orbitals (L=3) on the f-shell - fourteen electrons
Magnetic quantum number m l
There are three orbitals on the p-shell, they are denoted by numbers from -L to +L, that is, for the p-shell (L=1) there are orbitals "-1", "0" and "1". The magnetic quantum number is denoted by the letter m l .
Inside the shell, it is easier for electrons to be located in different orbitals, so the first electrons fill one for each orbital, and then its pair is added to each.
Consider a d-shell:
The d-shell corresponds to the value L=2, that is, five orbitals (-2,-1,0,1 and 2), the first five electrons fill the shell, taking the values M l =-2,M l =-1,M l =0 , M l =1, M l =2.
Spin quantum number m s
Spin is the direction of rotation of an electron around its axis, there are two directions, so the spin quantum number has two values: +1/2 and -1/2. Only two electrons with opposite spins can be on the same energy sublevel. The spin quantum number is denoted m s
Principal quantum number n
The main quantum number is the energy level, at the moment seven energy levels are known, each is denoted by an Arabic numeral: 1,2,3,...7. The number of shells at each level is equal to the level number: there is one shell on the first level, two on the second, and so on.
Electron number
So, any electron can be described by four quantum numbers, the combination of these numbers is unique for each position of the electron, let's take the first electron, the lowest energy level is N=1, one shell is located on the first level, the first shell at any level has the shape of a ball (s -shell), i.e. L=0, the magnetic quantum number can take only one value, M l =0 and the spin will be equal to +1/2. If we take the fifth electron (in whatever atom it is), then the main quantum numbers for it will be: N=2, L=1, M=-1, spin 1/2.
Since the nuclei of reacting atoms remain unchanged during chemical reactions (with the exception of radioactive transformations), the chemical properties of atoms depend on the structure of their electron shells. Theory electronic structure of the atom based on 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 division of 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 main quantum number n = 1, 2, 3,; orbital (azimuth) 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 that have the same set of four quantum numbers n,l,m l , m s; sets 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, moreover, in the energy layer with the given value n can be no more than 2n 2 electrons. Sets of electrons with the same quantum numbers n and l, form sublevels, denoted as they move away from the core as s, p, d, f.
The probabilistic finding of the position of an electron in the 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 the 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 the electron shells of atoms change, 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 substantiation of the periodic system of chemical elements includes data on the structure of the atom, confirming the existence of a relationship 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 the atomic nuclei, the number of electrons in the outer 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 the elements: at the beginning of the period (except for the first) there is alkali metal, then there is a gradual weakening of the metallic properties and an increase in non-metallic ones. This regularity can be traced for the elements of the second period in table 4.2.
Table 4.2.
In large periods, with an increase in the charge of nuclei, the filling of levels with electrons is more difficult, which explains the more complex change in the properties of elements compared to elements of small periods.
The same nature of the properties of chemical elements in subgroups is explained by the similar structure of the external energy level, as shown in tab. 4.3 illustrating the sequence of electron filling of energy levels for subgroups of alkali metals.
Table 4.3.
The group number, as a rule, 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 in the periodic table, the elements are not in ascending order of atomic masses: 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 the 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 system are filled in the interval from Z=1 before Z=114, however, the periodic table is not complete, the discovery of new transuranium elements is possible.
Electrons
The concept of an atom originated in the ancient world to denote the particles of matter. In Greek, atom means "indivisible".
The Irish physicist Stoney, on the basis of experiments, came to the conclusion that electricity is carried by the smallest particles that exist in the atoms of all chemical elements. In 1891, Stoney proposed to call these particles electrons, which in Greek means "amber". A few years after the electron got its name, English physicist Joseph Thomson and French physicist Jean Perrin proved that electrons carry a negative charge. This is the smallest negative charge, which in chemistry is taken as a unit (-1). Thomson even managed to determine the speed of the electron (the speed of an electron in orbit is inversely proportional to the orbit number n. The radii of the orbits grow in proportion to the square of the orbit number. In the first orbit of the hydrogen atom (n=1; Z=1), the speed is ≈ 2.2 106 m / c, that is, about a hundred times less than the speed of light c=3 108 m/s.) and the mass of an electron (it is almost 2000 times less than the mass of a hydrogen atom).
The state of electrons in an atom
The state of an electron in an atom is a set of information about the energy of a particular electron and the space in which it is located. An electron in an atom does not have a trajectory of motion, i.e., one can only speak of the probability of finding it in the space around the nucleus.
It can be located in any part of this space surrounding the nucleus, and the totality of its various positions is considered as an electron cloud with a certain negative charge density. Figuratively, this can be imagined 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 points. Overlaying countless such photographs would result in a picture of an electron cloud with the highest density where there will be most of these points.
The space around the atomic nucleus, in which the electron is most likely to be found, is called the orbital. It contains approximately 90% e-cloud, and this means that about 90% of the time the electron is in this part of space. Distinguished by shape 4 currently known types of orbitals, which are denoted by 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 electron 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 main 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 will be 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 largest number of electrons in the energy level is determined by the formula:
N = 2n2,
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.
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 sublevel; the second - two; third - three; fourth - four sublevels. Sublevels, in turn, are formed by orbitals. Each valuen corresponds to the number of orbitals equal to n.
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.
Protons and neutrons
An atom of any chemical element is comparable to a tiny solar system. Therefore, such a model of the atom, proposed by E. Rutherford, is called planetary.
The atomic nucleus, in which the entire mass of the atom is concentrated, consists of particles of two types - protons and neutrons.
Protons have a charge equal to the charge of electrons, but opposite in sign (+1), and a mass equal to the mass of a hydrogen atom (it is accepted in chemistry as a unit). Neutrons carry no charge, they are neutral and have a mass equal to that of a proton.
Protons and neutrons are collectively called nucleons (from the Latin nucleus - nucleus). The sum of the number of protons and neutrons in an atom is called the mass number. For example, the mass number of an aluminum atom:
13 + 14 = 27
number of protons 13, number of neutrons 14, mass number 27
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 represent e - .
Because the atom electrically neutral, it is also obvious that the number of protons and electrons in an atom is the same. It is equal to the serial number of the chemical element assigned to it in the Periodic system. The mass of an atom is made up 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, you can find the number of neutrons (N) using the formula:
N=A-Z
For example, the number of neutrons in an iron atom is:
56 — 26 = 30
isotopes
Varieties of atoms of the same element that have the same nuclear charge but different mass numbers are called isotopes. Chemical elements found in nature are a mixture of isotopes. So, carbon has three isotopes with a mass of 12, 13, 14; oxygen - three isotopes with a mass of 16, 17, 18, etc. The relative atomic mass of a chemical element usually given in the Periodic System is the average value of the atomic masses of a natural mixture of isotopes of a given element, taking into account their relative content in nature. The chemical properties of the isotopes of most chemical elements are exactly the same. However, hydrogen isotopes differ greatly in properties due to the dramatic fold increase in their relative atomic mass; they have even been given individual names and chemical symbols.
Elements of the first period
Scheme of the electronic structure of the hydrogen atom:
Schemes of the electronic structure of atoms show the distribution of electrons over electronic layers (energy levels).
The graphical electronic formula of the hydrogen atom (shows 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 orbits.
In a helium atom, the first electron layer is completed - it has 2 electrons. Hydrogen and helium are s-elements; for these atoms, the s-orbital is filled with electrons.
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 completed - it has 8 electrons.
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.
A 3s electron orbital is completed at the magnesium atom. Na and Mg are s-elements.
For aluminum and subsequent elements, the 3p sublevel is filled with electrons.
The elements of the third period have unfilled 3d orbitals.
All elements from Al to Ar are p-elements. s- and p-elements form the main subgroups in the Periodic system.
Elements of the fourth - seventh periods
A fourth electron layer appears at the potassium and calcium atoms, the 4s sublevel is filled, since it has less energy than the 3d sublevel.
K, Ca - s-elements included in the main subgroups. For atoms from Sc to Zn, the 3d sublevel is filled with electrons. These are 3d elements. They are included in 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 4s- 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 there can only be 32 electrons in the fourth electron layer; the 4d- and 4f-sublevels of the krypton atom still remain unfilled. The elements of the fifth period are filling the sub-levels in the following order: 5s - 4d - 5p. And there are also exceptions related to " failure» electrons, y 41 Nb, 42 Mo, 44 Ru, 45 Rh, 46 Pd, 47 Ag.
In the sixth and seventh periods, f-elements appear, i.e., elements in which the 4f- and 5f-sublevels of the third outer electronic layer are filled, respectively.
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 Cs and 56 Ba - 6s-elements; 57 La … 6s 2 5d x - 5d element; 58 Ce - 71 Lu - 4f elements; 72 Hf - 80 Hg - 5d elements; 81 T1 - 86 Rn - 6d elements. But even here 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 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.
- 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.
- 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 intercalary decades of large periods located between s- and p-elements. They are also called transition elements.
- f-elements. The f-sublevel of the third outside level of the atom is filled with electrons; these include the lanthanides and antinoids.
The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons having opposite (antiparallel) spins (translated from English - “spindle”), i.e. having such properties that can be conditionally imagined as the rotation of an electron around its imaginary axis: clockwise or counterclockwise.
This principle is called Pauli principle. If there is one electron in the orbital, then it is called unpaired, if there are two, then these are paired electrons, that is, electrons with opposite spins. The figure shows a diagram of the division of energy levels into sublevels and the order in which they are filled.
Very often, the structure of the 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 and F. Hund's rule, according to which electrons occupy free cells, first one at a time and at the same time have the same spin value, and only then pair, but the spins, according to the Pauli principle, will already be oppositely directed.
Hund's rule and Pauli's principle
Hund's rule- the rule of quantum chemistry, which determines the order of filling the orbitals of a certain sublayer and is formulated as follows: the total value of the spin quantum number of electrons of this sublayer should be maximum. Formulated by Friedrich Hund in 1925.
This means that in each of the orbitals of the sublayer, one electron is first filled, and only after the unfilled orbitals are exhausted, a second electron is added to this orbital. In this case, there are two electrons with half-integer spins of the opposite sign in one orbital, which pair (form a two-electron cloud) and, as a result, the total spin of the orbital becomes equal to zero.
Other wording: Below in energy lies the atomic term for which two conditions are satisfied.
- Multiplicity is maximum
- When the multiplicities coincide, the total orbital momentum L is maximum.
Let's analyze this rule using the example of filling the orbitals of the p-sublevel p- elements of the second period (that is, from boron to neon (in the diagram below, horizontal lines indicate orbitals, vertical arrows indicate electrons, and the direction of the arrow indicates the orientation of the spin).
Klechkovsky's rule
Klechkovsky's rule - as the total number of electrons in atoms increases (with an increase in the charges of their nuclei, or the ordinal numbers of chemical elements), atomic orbitals are populated in such a way that the appearance of electrons in higher-energy orbitals depends only on the principal quantum number n and does not depend on all other quantum numbers. numbers, including those from l. Physically, this means that in a hydrogen-like atom (in the absence of interelectron repulsion) the orbital energy of an electron is determined only by the spatial remoteness of the electron charge density from the nucleus and does not depend on the features of its motion in the field of the nucleus.
Klechkovsky's empirical rule and the sequence of sequences of a somewhat contradictory real energy sequence of atomic orbitals arising from it only in two cases of the same type: for atoms Cr, Cu, Nb, Mo, Ru, Rh, Pd, Ag, Pt, Au, there is a “failure” of an electron with s - sublevel of the outer layer to the d-sublevel of the previous layer, which leads to an energetically more stable state of the atom, namely: after filling the orbital 6 with two electrons s
