Periodic law development can be divided into 3 stages

Topic: Chemical elements.


Plan:

1. 
   Periodic law development can be divided into 3 stages
   
2. 
   Electronic analogy. Kaynosymmetry
   
3. 
   Periodic properties of elements
   
4. 
   The dependence of the elements of the iv group on ionization
   

Specificity and formation of cycles in the formation of atomic shells with electrons s, p-, d-, f- elements and their place in the periodic table. Groups, Periods, Main and Lateral Subgroups. The boundaries of the periodic table.
Periodicity of atomic properties. Orbital and effective radii. Van der Waals, metallic and ionic radii. Changes in atomic and ionic radii over periods and groups are effects of s- and p-compression. Kainosymmetric elements.
Various studies have been conducted from Kadima to systematize the chemical elements. (Examples of I. Debereiner, A. Shankurtua, Ch. Odling, Dj. Newlendes and others). In 1829, I. Debereiner was the first to study the relationship between the properties of chemical elements and the atomic weights of the elements, and created the law of placing many similar elements into groups of three, which consisted of 21 elements:

1. 
   cI , Br, I;
   
2. 
   S, Se, Te ;
   
3. 
   Li, Na, K;
   
4. 
   Ca, Si, Ba;
   
5. 
   Fe, Co, Ni;
   
6. 
   Os , Ir , Pt;
   
7. 
   

In 1862 esa French chemists A.Shankurtua chemical atomic masses of elements spiralsimon in order increasing to go according to posted . Here _ properties to each other similar which was elements groups yield to be observed . In 1857, Ch. Odling created a system of 57 elements in ascending order of atomic masses. In 1866, the octave law was created by J. Newlendes. He examined the relationship between the properties of chemical elements and atomic weights and found that there were similarities between each of the 8 elements, and created the following table:
H - 1 F - 8 cI - 15
Li - 2 Na - 9 K - 16
Be - 3 Mg - 10 Ca - 17
B - 4 Al - 11 Ti - 18
C - 7 Si - 12 Cr - 19
N - 6 P - 13 Mn - 20
O - 7 S - 14 Te - 21
Then, in 1864, German scientists Odling and Lothar-Meyers also tried to create a periodic table, but they also could not draw a definite scientific conclusion from their research.
In 1869, the periodic table of elements, created by DIMendeleev on the basis of the periodic law, became of universal importance. When DIMendeleev created the periodic table of elements, 63 elements were originally known. he placed them in a system that formed the word depending on the high oxidation states of the elements, predicting the properties of the unknown elements and finding them during life. These elements were discovered within 10 years of DIMendeleev's lifetime, namely, the discovery of aluminum (gallium) by Locock de Buabodran in 1875, ecabor (scandium) by the Swiss scientist Nilson Kleve in 1879, and ecacilic (Germanium) by Winkler in 1885. their very close proximity to the properties predicted by DIMendeleev showed that DIMendeleev further constructed the system correctly.
The periodic law, discovered by Dimendeleev in 1869, is one of the most fundamental laws in modern natural science. It is of great importance not only in chemistry but in the whole natural sciences as it signifies the material unity of the world. und the essence of chemistry as a science, that is, under the influence of quantitative changes in composition, are qualitative changes. The role of periodic law in the development of other natural sciences, physics, geochemistry, cosmochemistry is also great. Its importance is that elements are not classified by a single atomic weight. It predicts the properties of each element depending on its location in the system. This applies not only to the physical properties of ordinary substances, but also to the whole chemical properties. It allows to know the interaction, structure, formation, composition and properties of binary and more complex substances, acid-base, oxidation-reduction and other properties of elements. Using the periodic law, Mendeleev predicted the properties of elements that were still unknown. It is well known that the power of a true scientific theory lies not only in explaining the facts obtained in that sauce, but also in being able to see new facts. Describing the properties of the elements in the same order implies that in the periodic table, each element is precise, fixed, and constant. This is called the interval (position) of the positions (state). It is known that in the DIMendeleev system the position of elements is determined not only by its sequence number, but also by the number (series) and group of the period in which it stands.
Even in the form of the most common modern periodic table, the variant mode order of an element is not always preserved. Therefore, a general criterion that uniformly determines the position (position) of an element is necessary. Mendeleev himself chose the chemical properties of the elements as such a criterion. He considered its chemical properties to be the main characteristic in relation to the value of atomic masses. Therefore, he changed the positions of the elements (18Ar - 19K, 27Co - 28Ni, 52Te - 53I), that is, showed that the similarity in the groups exhibits chemical properties. Later, different scientists proposed different variants of the system, which were based on different, in some cases specific criteria, and now there are more than 400 variants of the system. Based on the development of the electronic theory of atomic structure, it was determined that the chemical properties of elements are functions of their atomic electronic structures. On this basis, it has been proved that it is expedient to select the electronic structure of the atom as an objective criterion for determining the state of the element in the periodic table.
Periodic law development can be divided into 3 stages.
In the first stage, the atomic mass is chosen as the main argument for determining the properties of the elements, and Mendeleev's periodic law is described on this basis as follows:
"The properties of simple substances, as well as the shape and properties of elemental compounds, depend periodically on the increasing atomic weights of the elements."
The second stage proved to determine the atomic number - the charge of the atomic nucleus. The discovery of isotopes and isobars has shown that the real argument that determines the nature of an element is its nuclear charge, not its atomic mass.
It is true that an atomic mass of isobars (40Ar, 40K, 40Ca) belongs to atoms of different elements, and that nuclear charges belong to the same element, even though the atomic masses of the same atoms - isotopes (160, 170, 180) are different. proved. Therefore, the periodic law was redefined:
"The shape and properties of simple substances, as well as elements, depend periodically on the charge of the nuclei of their atoms."
This change is of a fundamental nature and indicates a new qualitative level in the understanding of the nature of the element. the reason for the change was unclear.
Only in the third stage, based on the development of the quantum mechanical theory of the electronic structure, the physical meaning of the periodic law was revealed.
The essence of periodicity is based on the periodic repetition of similar valence electron configurations at high energy levels and the presence of the relative capacitance of the electron shells.
The structure and development of the periodic table were carried out in the following order.
The periodic table of elements includes 7 periods, 8 rups, and 10 rows.
The horizontal series of elements starting with alkali metals and ending with inert gases is called a period. Periods indicate the number of electronic layers of elements.
Periods I, II, III consist of one row and are called minor periods. Periods IV, V, and VI contain two rows, called major periods. Period VII is called the unfinished period, as it begins with alkali metals and does not reach inert gases.
A group of elements whose chemical properties are similar and have the same number of electrons in the outer electron shell is called a group. In which group the element is located, its highest oxidation state is equal to the group number, ie the number of electrons in the outer electron shell is equal to the number of that guru. The groups are divided into the main group (main) and additional (lateral). The main group elements are only s and p elements, while the additional group elements are d and f - elements. In addition to the electrons in the outer electron shell, the electrons in the outer inner electron shell are also valence electrons. So they are different from each other.
Periodic law is one of the basic laws of nature, which indicates the unity of the quantitative (nuclear charge, number of electrons and atomic masses) and qualitative (distribution of electrons, set of properties) elements.
The structure of the atom is based on modern ideas, the affiliation of the element to a specific period is determined by the number of electron layers in the normal, unexcited state of the atom.
The dvr number is equal to the outer electron layer number, which is filled with unfinished electrons.
The affiliation of an element to a particular group is determined by the total number of valence electrons in the outer and outer inner layers.
For example : 24Cr - [ Ar ] 183d 54s 1 and 16S - [Ne] 103s 23p 4.
Sixth group elements both atoms are 6 in number valent electrons has _ periods and to groups division by Mendeleev included of the element clear group affiliation his chemical property , high valence oxide and form of hydroxides and character based on detected . Really to each other similar not metallic chrome and metallmas oltingugurt group number right coming high oxidation level a kind of containing CrO3, SO3 oxides yield makes their _ properties are also similar ( acidic ). To them right coming hydroxides bright expressed acid property has H2 CrO4 - chromate and H2 SO4 - sulfate acids . That's right so periodic _ system in groups types depending on not without floors fill possible which was a kind of number electrons which was elements combines _ That's right so merge many similar (analog) species arrange opportunity gives _ Of the elements such general similarity species group analogues ( similar groups ) _ and they are group number appropriate only high oxidation levels show will be . For this sign, the main and auxiliary groups (group A and group V) are combined into one group.
Group III - B, Al, Ga, In, Tl (ns2np1) and the scandium subgroup [ns2 (n-1) d1], ie elements with the same valence electrons (3). A similar situation is characteristic of other guppies of the system.
Group similarity does not mean all the specific properties of the elements included in this group, as it often occurs as a common sign, based on the number of valence electrons, regardless of the type of valence orbitals.
This similarity loses its strength at low oxidation states of the elements, especially in the free state. But in each group it is possible to distinguish elements that show a deep resemblance to each other. Such similarity is manifested not only at high oxidation states, but at all intermediate oxidation states, only in the same valence electrons, in orbitals of the same type where the electrons are located. Based on this symbol, it is divided into subgroups in one group. The elements in a subgroup have very close similarities in properties, based on the fact that they have the same type of valence orbitals charged by electrons. Such an analogy, which has a much deeper resemblance, is called a typical resemblance.
Thus, in a subgroup - a type (analog type) similar to the elements belonging to this group, which have the same type of valence orbitals.
For example, the type III mentioned above is called a group-like type because all elements have the same valence electron orbitals (ns2 np1). The elements of the scandium subgroup also form additional III B - group - similar types, because for these too ns2 (n-1) d1 is the same.
From the electronic point of view of atomic structure, the period number has a bright physical meaning. It corresponds to the main quantum numeric value and is characterized by complementary or completed s and p layers.
Each phase valence electron configuration begins with ns1 and ends with ns2np6 stable configuration.
Periods in which atoms are filled only with s and p layers are called small periods. The first of these is the 3 period examples. 2, 8, 8 elements correspond to them. The number of elements in the first and second periods corresponds to the maximum capacity of the electric layers ( p = 1, p = 2). The capacitance of the third electron layer is higher (more) than the number of electrons. This is due to the presence of a 3d-orbital, which is empty, filling it with electrons is only an energetic tin in the fourth period. Thus, from the fourth period onwards, the order in which the electron shells are filled with electrons is disturbed, and ns, np are in the range of elements. The d-element decadence appears, in which the outer (n-1) d layer is filled from the outside. Such a structure will have the fourth and fifth periods, and they will have 18 elements. in the sixth and seventh periods, in addition to the decade of d-elements, there is a family of elements in which the (n-2) f-layer is paid. These periods consisted of 32 elements.
In addition to the s and p elements, periods consisting of the d-element decade and the f-element family are called major periods.
In addition to the outer layers of s and p, atoms in large cycles also have inner (n-1) d and (n-2) f-layers, which are also valence electrons.
The chemical properties of elements are determined by the specific properties of the filling of known atomic orbitals. Therefore, in small periods, there are a total of 8 elements in a row, the outermost layers of which are filled with electrons in their atoms, and the properties of the element change dramatically from the alkali metal to the inert gas as they move from one to the other.
In large periods, the s and p elements break down according to this law, while in the d-elements, the outer layer remains unchanged (ns2), and the properties of the outer second layer are filled with electrons, so their properties change more evenly. All d-elements are metals. To a greater extent , this feature is characteristic of f-elements, as they fill the third floor from the outside to the inside. The chemical properties of all f-elements are close to each other.


ELECTRONIC ANALOGY. KAYNOSYMMETRY
From the point of view of electronic structure (2-3), the elements of small circuits are characterized by their unique properties in comparison with all other elements.
This peculiarity is that the elements of the second and third periods do not fill the inner d- and f-layers, under the layer of valence electrons, the previous period is the inert gas atomic layer (ostov). A similar phenomenon is observed in large periods - in the elements. For example, the electronic configurations of the Na and Cs atoms [Xe] 54S1 and [Ne] 106S1, the configurations of the Na + and Cs + ions are similar to those of the electronic structures of ye and Ne Xe. The p elements in small periods are also similar to the s elements. But large periods are different in p-elements. They have valence electrons p s. After the p p layers, the ( p - 1) d layer is filled over the inert gas layer . for example: S and Se electronic configurations:
16 S = [Ne] 103s23p 4, 34Se = [ Ar ] 183d104s24p4
The formal oxidation states are S + 6 and Se + 6, and the ions are: S + 6 = [Ne] 10 Se + 6 = [Ar] 183d10. Thus, it is difficult to expect complete similarity of properties of S + 6 and Se + 6 at high oxidation states. However, at the + and lower -2 oxidation states of S and Se, complete electronic analogies, ie similar valence electron configurations, are observed. On the other hand, there is no similarity between the elements of group VI A and VI B (Example: S and Cr) in the intermediate oxidation states.
6 S = 1s22s22p63s23p 4; 24Cr = 1s22 s 22p63 s 23p64 s 13d5
But S and Cr are high oxidation levels similarity to the surface because inert gas in S + 6 and Cr + 6 is S + 6 [Ne] 10, Cr + 6 [ Ar ] 18, electron to the floor teng will be . Electron structure that's it similar legitimacy the rest are also observed in groups . Valent electron configurations similar which was to the elements is called an electronic analog . That's right so small _ period elements head gurappa elements and additional gruppacha elements ( each kind of oxidation levels ) electrons analogy show does and all group image determines _ Therefore for small period elements typical elements is called . This term D. I. Mendeleev by included . That's right so complete _ and incomplete electronic analog concept critique possible . all oxidation levels similar electron structure has which was to the elements full analog elements is called . This is theirs chemical properties according to close in similarities determined . For example : Group VI elements between complete electron analog: oxygen and oltingugurt 80 = [He] 2 2S2 2P4, 16 S = [Ne] 3S23P4, selenium , tellurium and polonium :
34 Se = [ Ar ] 183d104S24P4,
52 Te = [Kr] 364d105S25P4,
84 Po = [ Xe ] 544f145d106S26P4: and chrome molybdenum and tungsten : 24 Cr = [ Ar ] 183d54S1,
42 Mo = [Kr] 364d55S1,
74 W = [ Xe ] 544f145d46S2.
Polish and in tungsten other from the elements different completed 4f floor available , but it is much pit located for properties kam ta ' sir reaches and electron analog character does not break .
Typical elements atoms-oxygen and oltingugurt electron structures according to selen subgroups higher than the elements oxidation levels chrome subgroups from the elements from above from all oxidation levels difference does . This means that oxygen and sulfur are incomplete analogs to other elements of group VI. At the same time, the similarity in the electronic structure between the typical elements and the elements of the selenium subgroup is much closer. They can show not only group similarity, but also typical similarity, as noted above. The electronic similarity in group VI can be shown in the following diagram.
Incomplete electronic analog, analog only at high oxidation states. Complete and incomplete electronic similarity was first introduced by BVNekrasov. This means that if all oxidation states have the same electronic configuration, we call them complete electronic analogues, if they have the same electronic oxidation configuration, they are called incomplete electronic analogues.
For all groups in the same periodic table, it is possible to show an electronic analogy other than (VIII) gr. This is due to the similarity of the structures of the ns 1-2 electron orbitals of all the elements in a group. These elements do not have an inner d-layer on the outside. The typical elements of the two groups are lithium, sodium and beryllium, the auxiliary group elements of magnesium have incomplete electronic analogs with copper, silver, gold and zinc, cadmium, mercury. these elements belong to the group of typical analogs and are fully electronic analogues with each other. There is no complete analogy in groups III and VII (through). Each group is divided into 3 families of fully electronic analogs: 1. Typical elements; 2. The remaining elements of the other group; 3. Auxiliary subgroup elements. These families are connected to each other by an incomplete analogy of typical elements. Since the elements in the main subgroups have typical analogs, the electronic analogy is much more freely expressed. The difference in the electronic analogy of these groups with respect to groups I and II is the appearance of the d-layer. In terms of the electronic structure of the typical elements of the second and third periods, there is a vacant 3d-layer, which can participate in chemical reactions under certain conditions. this occurs when the valence electrons in the 3s and 3p layers are transferred to the same 3d layer by the principal quantum number. In the elements of the second period, the transfer to the 3s vacant layer is not energetically convenient: that is, the energy expended in the transfer to the vacant 3d layer cannot be compensated by the energy released in the formation of this garden additional garden. For this reason alone, the fifth valence state of nitrogen does not occur, and for phosphorus, this condition is normal. PS15s are known to be NG3 for nitrogen. Indeed, the structure of the valence orbitals of nitrogen ( p = 2) and phosphorus can be shown as follows:
From the point of view of the spin theory of valence, the covalence of nitrogen is equal to 3 and that of phosphorus is equal to 3 (3 current electrons). For example: in ammonium and phosphonium cations:
However, unlike nitrogen, phosphorus can be converted to a 5-covalent state.
The 5-valent compounds of phosphorus are thus explained. The peculiarity of the typical elements of the first series, on the other hand, is that in them the p-orbitals first formed. In these orbitals, the radial distribution function of the electron density has a single maximum. the orbitals that occur for the first time are called kainosymmetry. (Kaynos-Greek-new, i.e. new type symmetries of orbitals). In all other orbitals there will be additional maxima on the radial distribution curves of the electron density.
Thus, kainosymmetric orbitals are characterized by the absence of internally filled orbitals in such symmetry. This enhances the binding of kainosymmetric electrons to the nucleus, increases the radii of the atomic orbitals, increases the ionization potentials, and leads to a weakening of the metallic properties of the kainosymmetric elements relative to the non-kainosymmetric elements.
Periodicity in the chemical properties of simple and complex compounds of elements. The law of change of chemical activity of metals and non-metals in periods and groups. Changes in the acid-base properties of oxides and hydroxides in periods and groups. Classical internal and secondary periodicity.
The complexity of the Schrödinger equation does not allow an accurate calculation of the energy of multi-electron orbitals. therefore, several methods are used to calculate nuclear energy. One of the most convenient and universal methods for determining the energy of atomic orbitals experimentally is the spectroscopic analysis method. As we know, atomic spectra form a set of spectral lines (sovokupnost), that is, each series determines the transfer of electrons from different orbitals to another nucleus. The energy released when an electron passes from one orbital to another is called the ionization energy or ionization potential.
Ye is the base - K = I, where I is the ionization energy, Ye is the energy of the electron that makes up the atom, and Ye is the base energy of the electron.
If the frequency of electromagnetic oscillations from a small wave field yi and the ionization energy Ii can be expressed by Planck's formula. Ii = hyi.
If there is a multi-electron atom, it is convenient to determine the breakdown energy of the electron in other ways, that is, using the photoionization method, and so on.
Ionization energy is one of the main characteristics of a chemical element, characterizing the chemical properties of the element by showing how it holds the electron in its atomic orbital. so less energy is required to cut off the first electron, while the latter have to expend more energy.
The elements with the lowest ionization energy are alkali metals because they have the lowest energy value of one electron in the outer layer around the atomic nucleus and easily transfer to the atom of another element, the series VIII - group elements with the highest ionization energy (inert gases). A schematic representation of the increase in ionization potential is shown in Figure 1.
it can be seen that the ionization energy increases from left to right in the (main) periods, and decreases in groups from top to bottom.
Binding energy - Ye. The energy that binds to the energy released or absorbed when atomic nuclei combine to form a neutral substance is called energy. Ye - kcal / mol, kal / mol, kdj / mol are expressed in units. The numerical value of the binding energy indicates how stable or unstable the same substance is.
Bond length - 1 - The distance between interconnected atomic nuclei is called the bond length. l = A is expressed in angstroms. 1A-10-8 cm ga teng.
The binding angle - l The angle between the interconnected atomic nuclei is called the binding angle. The contact angle is measured in l - degrees and minutes.
Indicates the geometric shape (structure) of a substance with a size such as bond length and bond angle .
For example: The length of the bond between the H and O atoms in a water molecule 0,96 Ais 104.50. hence, the structure of the water molecule has a triangular shape.
Ionization potential - I The ability of an element to react can be reduced by the ionization potential - I and the electron affinity - Ye, ie,
a) The energy required to cut one electron from a normal nucleus is called the ionization potential of the same element. Its axis is I = kdj / atom or Ev \ atom E + I = E + Ie I1 The numerical value of the ionization potential indicates the metallic or non-metallic activity of the same element. The elements with the smallest ionization potential are alkaline and alkaline - earth metals. In the periodic table of elements, the ionization potentials of the elements increase as they move from left to right in periods. In groups, the ionization potentials of the elements decrease from top to bottom. If we study that the ionization potential of an element depends on the sequence number, then it is characterized by periodicity - extremum. The maximum for inert gases and the minimum for alkali metals. The reason for the maximum potential in inert gases is the stability of the electronic configurations p S2 p P6. The reason for its low is the effective shielding of the nuclear charge by the electrons in the outer p S orbital. The shielding effect is said to weaken the effect of this electron on the positive charge of the nucleus when another electron falls between them. Due to this, the electron interacts with Z eff (Z eff ) rather than with 2 . Slater says that the effective nuclear charge of the nucleus is Z eff = ZS, which is called the shielding constant of the S-nucleus using all the remaining electrons. Slater showed up to figure it out. The ionization potential varies between 5-15 B. It can be graphically described as follows.
Electron propensity - Ye The energy released when a single electron joins a normal atom is called the electron propensity of the same element.
Changes in the values of ionization potentials of elements in periods and groups.
From time to time the number of shots increases.
(Ye) E + e6E + Ye, Ye = kdj / mol or Home / atom.
The greater the electron affinity of an element, the stronger the non-metallic property of the same element. Accordingly, in the periodic table of elements, the electron affinity of the elements increases from left to right in periods, and in groups the electron affinity decreases from top to bottom. It can be represented graphically as follows:
The propensity for electrons is also a periodic variable property, which occurs in the lateral group of group VII of the periodic table, and at least in inert and alkaline earth metals.
Electromagnetism - EM. A chemical bond is a quantity that indicates the ability of an electron to go or bind when it is formed.
Malliken explained electromagnetism as the sum of the ionization potential + electron affinity.
I + E = Electromagnetism.
This includes the Malliken and Poling scales. On the Pauling scale, the relative electronegativity of the lithium atom was assumed to be 1.
In 1927, Poling introduced the concept of relative electronegativity (NEM) values to explain the metallic and non-metallic properties of elements in the periodic table of elements. The NEMs of the elements are given in the form of a scale in the table, which is called the Pauling scale. It is mainly used in the study of chemical bonds. The NEMs of the elements increase from left to right in periods (for the main group elements), and in groups the value of NEMs decreases from top to bottom. It can be displayed graphically as follows:
Li - -> Be - -> B - -> C - -> N - -> O - -> F.
1.0 1.5 2.0 2.5 3.0 3.6 4.1
Na 0, 97
K 0.85
Rb 0, 86
Cs 0, 86
The general electronic configuration of the elements is represented by the following formula:
[E] (n-2) f (n-1) d ns np
where [E] is the electronic configuration of inert gases a, b, g, s, f, d, s, p, the number of electrons in the electron shell:
If a = 0, b = 0, s = 0, the general electronic configuration is characteristic of us elements when [E] p S b. These are the main subgroup elements of groups 1 and 2. Their properties are unique. For example: acid-base properties of protons in solution, hypersensitivity of liquid helium.
If a = 0 and = 0, the electronic configurations of these atoms are [E] p sa p pb, which are sp elements. They include elements of the main group of groups III-VIII. This can be studied in two groups p s2 p p6 and p s2 p p1 to p s2 p p5. The first is inert gases, which are said to be stable and inactive. In 1962, Bartlett obtained more than 100 compounds of krypton, xenon, and radon. These were the basis for placing them in the main group of Group VIII. The elements in the second group are both metals and non-metals, which can give or receive electrons to fill their electron shells to the state of inert gases.
If a = 0, s = 0, the electronic configuration is [E] ( p -1) db p sa, which are the elements sd, the elements between periods 4, 5, 6, and 7. These elements ( p -1) are typical metals characterized by the filling of d layers. Their properties do not differ much from each other, only at high oxidation states.
If s = 0, the electronic configuration is E ( p -2) fa ( p -1) db nsa. These are the sf and sdf elements and the typical metals are lanthanides and actinoids. Thus, the electronic configurations of the atoms show that the properties of the elements depend periodically on the increase in the sequence numbers.
Periodic properties of elements.
Periodicity is observed by studying the place in the system of the physical and chemical properties of the elements in the periodic table.
It was proved that as a result of periodic increase in the number of elements, their properties change periodically. The development of theories of atomic structure explained its physical nature, as well as an increase in electronic configurations, an increase in the number of head quanta, and a periodic change. Since the properties of elements depend on the electronic configuration of their atoms, it is felt that there is a periodicity. If we look at the property number diagram, we see a periodicity consisting of a minimum and a maximum. Periodicity can be seen in the simple change in the properties of free atoms, and in the complex in the formation of chemical compounds from particles in the gaseous state, in solution. As a classic example of periodicity, it can be said that the volume of atoms of simple substances, ionic radii, oxidation numbers, periodically depend on the sequence number. properties related to electronic configurations — ionization potential, electron susceptibility, electronegativity — are also periodic.
In 1915, Biron introduced to science the possibility of a monotonous change of properties in periods, that is, a secondary periodicity. The figure shows 4 groups of lateral monotonous and non-monotonous changes. The reason for the monotony is the increase in the effective charge of the germanium element due to the display of 4S electrons from 3d electrons.


SI - IONIZATION DEPENDENCE OF GROUP IV ELEMENTS
In conclusion, the laws of chemical activity of metals and non-metals in the periodic table of elements, as well as changes in the acid-base properties of oxides and hydroxides in periods and groups can be shown in the following scheme:
Periods from left to right in periods
oxidizing properties increase;
increased susceptibility to electrons;
non-metallic properties increase;
relative electromagnetism increases;
the acid strength of oxides and hydroxides increases, the basic property decreases.

• 
  atomic and ionic radii increase;
  
• 
  ionization potentials decreased ;
  
• 
  electrons inclination decreased ;
  
• 
  relative electronegativity decreased ;
  
• 
  return property increases ;
  
• 
  the basic properties of oxides and hydroxides increase;
  
• 
  in the water solubility increases ;
  

Atomic and ionic radii decrease;
The radius of the negative ion increases;
The metallic inclination increases;
Ionization potentials increase;
The metallic properties decrease and the non-metallic properties increase;
The acidity of oxides and hydroxides increases;
Oxidizing property increases .