Old wording. The discovery of the periodic law of chemical elements by D.I. Mendeleev. Basic principles of constructing the periodic system

: as the famous Russian chemist N.D. Zelinsky figuratively noted, Periodic law was "the discovery of the mutual connection of all atoms in the universe."

History

The search for the basis of natural classification and systematization chemical elements began long before the discovery of the Periodic Law. The difficulties encountered by natural scientists who were the first to work in this field were caused by the lack of experimental data: at the beginning of the 19th century, the number of known chemical elements was small, and the accepted values ​​of the atomic masses of many elements are incorrect.

Döbereiner's triads and the first systems of elements

In the early 60s of the XIX century, several works appeared at once that immediately preceded the Periodic Law.

Spiral de Chancourtois

Octaves of Newlands

Newlands Table (1866)

Soon after de Chancourtois spiral, the English scientist John Newlands made an attempt to compare the chemical properties of elements with their atomic masses. Arranging the elements in ascending order of their atomic masses, Newlands noticed that similarities in properties appear between every eighth element. The found pattern Newlands called the law of octaves by analogy with seven intervals of a musical scale. In his table, he arranged the chemical elements into vertical groups of seven elements each, and at the same time found that (with a slight change in the order of some elements) elements with similar chemical properties appear on the same horizontal line.

John Newlands, undoubtedly, was the first to give a number of elements arranged in ascending order of atomic masses, assigned a corresponding serial number to chemical elements, and noticed a systematic relationship between this order and the physicochemical properties of the elements. He wrote that in such a sequence the properties of elements are repeated, the equivalent weights (masses) of which differ by 7 units, or by a value that is a multiple of 7, that is, as if the eighth element in order repeats the properties of the first, as in music the eighth note repeats first. Newlands tried to make this dependence, which is indeed the case for the light elements, a universal character. In his table, similar elements were located in horizontal rows, however, elements with completely different properties were often in the same row. In addition, Newlands had to accommodate two elements in some cells; finally, the table did not contain empty spaces; as a result, the law of octaves was received with great skepticism.

Odling and Meier tables

Manifestations of the periodic law in relation to the electron affinity energy

The periodicity of the values ​​of the energies of the affinity of atoms for an electron is naturally explained by the same factors that have already been noted in the discussion of ionization potentials (see the definition of the energy of affinity for an electron).

The greatest affinity for the electron is possessed by p- elements of the VII group. The smallest electron affinity is for atoms with the s² (,,) and s²p 6 (,) configurations or with half-filled p-orbitals (,,):

Manifestations of the periodic law of electronegativity

Strictly speaking, an element cannot be ascribed permanent electronegativity. The electronegativity of an atom depends on many factors, in particular on the valence state of the atom, the formal oxidation state, the coordination number, the nature of the ligands that make up the environment of the atom in the molecular system, and some others. Recently, the so-called orbital electronegativity has been used more and more often to characterize electronegativity, which depends on the type of atomic orbital involved in the formation of a bond and on its electronic population, i.e. on whether the atomic orbital is occupied by a lone electron pair, is populated once by an unpaired electron, or is vacant. But, despite the well-known difficulties in the interpretation and definition of electronegativity, it always remains necessary for a qualitative description and prediction of the nature of bonds in a molecular system, including the binding energy, the distribution of the electronic charge and the degree of ionicity, the force constant, etc.

The periodicity of atomic electronegativity is an important component of the periodic law and can be easily explained based on the immutable, although not entirely unambiguous, dependence of the electronegativity values ​​on the corresponding values ​​of ionization energies and electron affinity.

In periods, there is a general tendency towards an increase in electronegativity, and in subgroups, its decline. The smallest electronegativity is for s-elements of group I, the highest is for p-elements of group VII.

Manifestations of the periodic law in relation to atomic and ionic radii

Rice. 4 Dependence of the orbital radii of atoms on the ordinal number of the element.

The periodic nature of changes in the size of atoms and ions has been known for a long time. The difficulty here is that, due to the wave nature of the electronic motion, atoms do not have strictly defined sizes. Since direct determination of the absolute sizes (radii) of isolated atoms is impossible, in this case their empirical values ​​are often used. They are obtained from the measured internuclear distances in crystals and free molecules, breaking each internuclear distance into two parts and equating one of them to the radius of the first (of two connected by the corresponding chemical bond) atom, and the other to the radius of the second atom. This division takes into account various factors, including the nature chemical bond, the oxidation state of two bound atoms, the nature of the coordination of each of them, etc. In this way, the so-called metallic, covalent, ionic and van der Waals radii are obtained. Van der Waals radii should be considered as the radii of unbound atoms; they are found by internuclear distances in solid or liquid substances, where atoms are in close proximity to each other (for example, atoms in solid argon or atoms from two neighboring N 2 molecules in solid nitrogen), but are not linked to each other by any chemical bond ...

But obviously the best description the effective size of an isolated atom is the theoretically calculated position (distance from the nucleus) of the main maximum of the charge density of its outer electrons. This is the so-called orbital radius of the atom. The periodicity in the change in the values ​​of the orbital atomic radii, depending on the ordinal number of the element, manifests itself quite clearly (see Fig. 4), and the main points here consist in the presence of very pronounced maxima corresponding to the atoms of alkali metals, and the same minima corresponding to noble gases ... The decrease in the values ​​of the orbital atomic radii during the transition from an alkali metal to the corresponding (nearest) noble gas is, with the exception of the series -, non-monotonic, especially when between alkali metal and a noble gas of the families of transition elements (metals) and lanthanides or actinides. In large periods in families d- and f- elements, a less sharp decrease in radii is observed, since the filling of the orbitals with electrons occurs in the pre-outer layer. In subgroups of elements, the radii of atoms and ions of the same type generally increase.

Manifestations of the periodic law in relation to atomization energy

It should be emphasized that the oxidation state of an element, being a formal characteristic, does not give an idea of ​​either the effective charges of the atoms of this element in the compound, or the valence of the atoms, although the oxidation state is often called the formal valence. Many elements are capable of exhibiting not one but several different oxidation states. For example, for chlorine, all oxidation states are known from −1 to +7, although even ones are very unstable, and for manganese, from +2 to +7. The highest values ​​of the oxidation state change periodically depending on the ordinal number of the element, but this periodicity has complex nature... In the simplest case, in the series of elements from an alkali metal to a noble gas, the highest oxidation state increases from +1 (F) to +8 (O 4). In other cases, the highest oxidation state of the noble gas is lower (+4 F 4) than for the preceding halogen (+7 O 4 -). Therefore, on the curve of the periodic dependence of the highest oxidation state on the ordinal number of the element, the maxima fall either on the noble gas or on the halogen preceding it (the minima are always on the alkali metal). An exception is the series -, in which high oxidation states are generally unknown neither for halogen (), nor for a noble gas (), and the greatest value the highest oxidation state is possessed by the middle member of the series - nitrogen; therefore, in the series - the change in the highest oxidation state turns out to be passing through a maximum. In the general case, the increase in the highest oxidation state in the series of elements from an alkali metal to a halogen or to a noble gas is by no means monotonic, mainly due to the manifestation of high oxidation states with transition metals. For example, an increase in the highest oxidation state in the series - from +1 to +8 is "complicated" by the fact that for molybdenum, technetium and ruthenium such high oxidation states as +6 (О 3), +7 (2 О 7), + 8 (O 4).

Manifestations of the periodic law in relation to oxidative potential

One of the very important characteristics of a simple substance is its oxidative potential, which reflects the fundamental ability of a simple substance to interact with aqueous solutions, as well as the oxidation-reduction properties shown by it. Change in oxidative potentials simple substances depending on the serial number of the element, it is also periodic. But it should be borne in mind that the oxidative potential of a simple substance is influenced by various factors, which sometimes need to be considered individually. Therefore, the periodicity in the change in oxidation potentials should be interpreted very carefully.

/ Na + (aq)	/ Mg 2+ (aq)	/ Al 3+ (aq)
2.71V	2.37V	1.66V
/ K + (aq)	/ Ca 2+ (aq)	/ Sc 3+ (aq)
2.93V	2.87V	2.08V

You can find some definite sequences in the change in the oxidative potentials of simple substances. In particular, in the series of metals, when passing from alkaline to the elements following it, the oxidation potentials decrease (+ (aq), etc. - hydrated cation):

This is easily explained by an increase in the ionization energy of atoms with an increase in the number of removed valence electrons. Therefore, on the curve of the dependence of the oxidation potentials of simple substances on the ordinal number of the element, there are maxima corresponding to alkali metals. But this is not the only reason for the change in the oxidative potentials of simple substances.

Internal and secondary periodicity

s- and R-elements

Above, general trends in the nature of changes in the values ​​of the ionization energy of atoms, the energy of the affinity of atoms to the electron, electronegativity, atomic and ionic radii, the energy of atomization of simple substances, the oxidation state, oxidation potentials of simple substances from the atomic number of an element are considered. With a deeper study of these trends, one can find that the patterns in the change in the properties of elements in periods and groups are much more complicated. In the nature of changes in the properties of elements by period, internal periodicity is manifested, and in the group - secondary periodicity (discovered by E.V. Biron in 1915).

So, when passing from an s-element of group I to R-element of group VIII on the curve of the ionization energy of atoms and the curve of changing their radii have internal maxima and minima (see Fig. 1, 2, 4).

This indicates the internally periodic nature of the change in these properties over the period. The above regularities can be explained using the concept of nuclear screening.

The shielding effect of the nucleus is caused by the electrons of the inner layers, which, by shielding the nucleus, weaken the attraction of the outer electron to it. So, when going from beryllium 4 to boron 5, despite the increase in the nuclear charge, the ionization energy of atoms decreases:

Rice. 5 Diagram of the structure of the last levels of beryllium, 9.32 eV (left) and boron, 8.29 eV (right)

This is because the attraction to the core 2p-electron of boron atom is weakened due to the shielding action 2s-electrons.

It is clear that the screening of the nucleus increases with the number of internal electronic layers... Therefore, in the subgroups s- and R-elements, there is a tendency to a decrease in the ionization energy of atoms (see Fig. 1).

The decrease in the ionization energy from nitrogen 7 N to oxygen 8 O (see Fig. 1) is explained by the mutual repulsion of two electrons of the same orbital:

Rice. 6 Scheme of the structure of the last levels of nitrogen, 14.53 eV (left) and oxygen, 13.62 eV (right)

The effect of screening and mutual repulsion of electrons of one orbital also explains the internally periodic nature of the change over the period of atomic radii (see Fig. 4).

Rice. 7 Secondary periodic dependence of the radii of the atoms of the outer p-orbitals on the atomic number

Rice. 8 Secondary periodic dependence of the first ionization energy of atoms on the atomic number

Rice. 9 Radial distribution of electron density in the sodium atom

In the nature of the change in properties s- and R-elements in subgroups, a secondary periodicity is clearly observed (Fig. 7). To explain it, the concept of the penetration of electrons to the nucleus is used. As shown in Figure 9, an electron of any orbital for a certain time is in a region close to the nucleus. In other words, the outer electrons penetrate to the nucleus through the layers of inner electrons. As can be seen from Figure 9, external 3 s-electron of a sodium atom has a very significant probability of being near the nucleus in the region of internal TO- and L-electronic layers.

The concentration of electron density (the degree of penetration of electrons) at the same principal quantum number is highest for s-electron, less - for R-electron, even less - for d-electron, etc. For example, for n = 3, the degree of penetration decreases in the sequence 3 s>3p>3d(see fig. 10).

Rice. 10 Radial distribution of the probability of finding an electron (electron density) at a distance r from the core

It is clear that the penetration effect increases the strength of the bond between the outer electrons and the nucleus. Due to deeper penetration s-electrons in to a greater extent shielding the core than R-electrons, and the latter are stronger than d-electrons, etc.

Using the concept of electron penetration to the nucleus, let us consider the nature of the change in the radius of the atoms of elements in the carbon subgroup. In the series - - - - there is a general tendency to increase the radius of the atom (see Fig. 4, 7). However, this increase is non-monotonic. On going from Si to Ge, the external R-electrons penetrate the screen out of ten 3 d-electrons and thereby strengthen the bond with the nucleus and compress the electron shell of the atom. Downsize 6 p-orbitals of Pb compared to 5 R-orbital Sn is due to the penetration of 6 p- electrons under the double screen ten 5 d-electrons and fourteen 4 f-electrons. This also explains the non-monotonicity in the change in the ionization energy of atoms in the C-Pb series and its greater value for Pb as compared to the Sn atom (see Fig. 1).

d-Elements

In the outer layer of atoms d-elements (except for) there are 1-2 electrons ( ns-condition). The rest of the valence electrons are located in (n-1) d-state, that is, in the pre-outer layer.

Similar structure electronic shells atoms determines some general properties d-elements. Thus, their atoms are characterized by relatively low values ​​of the first ionization energy. As can be seen in Fig. 1, the nature of the change in the ionization energy of atoms over the period in the series d-elements are smoother than in a row s- and p-elements. When moving from d-element of the III group to d-element of the II group, the values ​​of the ionization energy change non-monotonically. Thus, in the segment of the curve (Fig. 1), two areas are visible, corresponding to the ionization energy of atoms, in which 3 d-orbitals one and two electrons each. Filling 3 d-orbitals, one electron each ends at (3d 5 4s 2), which is marked by a slight increase in the relative stability of the 4s 2 -configuration due to the penetration of 4s 2 -electrons under the shield of the 3d 5 -configuration. The highest value of the ionization energy has (3d 10 4s 2), which is in accordance with the complete completion of 3 d-sublayer and stabilization of the electron pair due to penetration under the screen 3 d 10 configurations.

In subgroups d-elements, the values ​​of the ionization energy of atoms generally increase. This can be explained by the effect of the penetration of electrons to the nucleus. So if u d- elements of the 4th period external 4 s-electrons penetrate the screen 3 d-electrons, then the elements of the 6th period have external 6 s-electrons already penetrate under the double screen 5 d- and 4 f-electrons. For example:

22 Ti ... 3d 2 4s 2	I = 6.82 eV
40 Zr… 3d 10 4s 2 4p 6 4d 2 5s 2	I = 6.84 eV
72 Hf… 4d 10 4f 14 5s 2 5p 6 5d 2 6s 2	I = 7.5 eV

Therefore, d-elements of the 6th period external b s-electrons are more firmly bound to the nucleus and, therefore, the ionization energy of atoms is greater than that of d-elements of the 4th period.

Atom sizes d-elements are intermediate between the sizes of atoms s- and p-elements of this period. The change in the radii of their atoms over the period is smoother than for s- and p-elements.

In subgroups d-elements, the radii of atoms generally increase. It is important to note the following feature: an increase in atomic and ionic radii in subgroups d-elements basically corresponds to the transition from the 4th element to the 5th period element. The corresponding radii of atoms d-elements of the 5th and 6th periods of this subgroup are approximately the same. This is explained by the fact that an increase in the radii due to an increase in the number of electron layers during the transition from the 5th to the 6th period is compensated by f-compression caused by filling with electrons 4 f-sublayer at f-elements of the 6th period. In this case f-compression is called lanthanoid... With similar electronic configurations of the outer layers and approximately the same sizes of atoms and ions for d-elements of the 5th and 6th periods of this subgroup are characterized by a special similarity of properties.

The elements of the scandium subgroup do not obey these patterns. For this subgroup, typical patterns are typical for neighboring subgroups s-elements.

The periodic law is the basis of chemical taxonomy

see also

Notes (edit)

Literature

1. Akhmetov N.S. Topical issues of the course inorganic chemistry... - M .: Education, 1991 .-- 224 p. - ISBN 5-09-002630-0
2. D. V. Korolkov Fundamentals of Inorganic Chemistry. - M .: Education, 1982 .-- 271 p.
3. Mendeleev D.I. Fundamentals of chemistry, vol. 2. M .: Goskhimizdat, 1947.389 p.
4. Mendeleev D.I.// Encyclopedic Dictionary of Brockhaus and Efron: In 86 volumes (82 volumes and 4 additional). - SPb. , 1890-1907.

First option Periodic Table of Elements was published by Dmitry Ivanovich Mendeleev in 1869 and was called "The experience of a system of elements."

DI. Mendeleev arranged 63 elements known at that time in the order of increasing atomic masses and obtained a natural series of chemical elements, in which he discovered periodic recurrence of chemical properties. This series of chemical elements is now known as the Periodic Law (formulated by D.I.Mendeleev):

The properties of simple bodies, as well as the shapes and properties of compounds of elements, are periodically dependent on the value of the atomic weights of the elements.

The modern wording of the law is as follows:

the properties of chemical elements, simple substances, as well as the composition and properties of compounds are periodically dependent on the values ​​of the charges of atomic nuclei.

Graphic image periodic law is the periodic table.

In the cell of each element, its most important characteristics are indicated.

Periodic table contains groups and periods.

Group- a column of the periodic system, in which chemical elements are located that have chemical similarity due to identical electronic configurations of the valence layer.

Periodic table of D.I. Mendeleev contains eight groups of elements. Each group consists of two subgroups: main (a) and secondary (b). The main subgroup contains s- and p- elements, in the side - d- elements.

Group names:

I-a Alkali metals.

II-a Alkaline earth metals.

V-a Pnictogens.

VI-a Chalcogenes.

VII-a Halogens.

VIII-a Noble (inert) gases.

Period is a sequence of elements, written as a string, arranged in the order of increasing charges of their nuclei. The period number corresponds to the number of electronic levels in the atom.

The period begins with an alkali metal (or hydrogen) and ends with a noble gas.

Parameter	Down the group	Period right
Core charge	Is increasing	Is increasing
Valence electrons	Does not change	Is increasing
Number energy levels	Is increasing	Does not change
Atom radius	Is increasing	Decreases
Electronegativity	Decreases	Is increasing
Metallic properties	Are increasing	Are decreasing
Oxidation state in higher oxide	Does not change	Is increasing
Oxidation state in hydrogen compounds (for elements of IV-VII groups)	Does not change	Is increasing

Modern periodic table of chemical elements of Mendeleev.

In 1871 Mendeleev's periodic law was formulated. By this time, 63 elements were known to science, and Dmitry Ivanovich Mendeleev ordered them on the basis of relative atomic mass. The modern periodic table has expanded significantly.

History

In 1869, while working on a chemistry textbook, Dmitry Mendeleev faced the problem of systematizing the material accumulated over many years by various scientists - his predecessors and contemporaries. Even before Mendeleev's work, attempts were made to systematize the elements, which served as prerequisites for the development of the periodic system.

Rice. 1. Mendeleev D.I ..

Element classification searches are summarized in the table.

Mendeleev ordered the elements according to their relative atomic mass, arranging them in ascending order. There are nineteen horizontal and six vertical rows in total. This was the first edition periodic table elements. This is the beginning of the history of the discovery of the periodic law.

It took the scientist almost three years to create a new, more perfect table. Six columns of elements turned into horizontal periods, each of which began with an alkali metal and ended with a non-metal (inert gases were not yet known). The horizontal rows formed eight vertical groups.

Unlike his colleagues, Mendeleev used two criteria for the distribution of elements:

• atomic mass;
• Chemical properties.

It turned out that there is a pattern between these two criteria. After a certain number of elements with increasing atomic mass, the properties begin to repeat themselves.

Rice. 2. The table compiled by Mendeleev.

Initially, the theory was not expressed mathematically and could not be fully confirmed experimentally. The physical meaning of the law became clear only after the creation of the atomic model. The point is to repeat the structure of the electron shells with a sequential increase in the charges of the nuclei, which is reflected in the chemical and physical properties elements.

Law

Having established the periodicity of changes in properties with an increase in atomic mass, Mendeleev in 1871 formulated a periodic law, which became fundamental in chemical science.

Dmitry Ivanovich determined that the properties of simple substances are periodically dependent on the relative atomic masses.

Science of the XIX century did not possess modern knowledge about the elements, therefore the modern formulation of the law is somewhat different from Mendeleev's. However, the essence remains the same.

WITH further development science, the structure of the atom was studied, which influenced the formulation of the periodic law. According to the modern periodic law, the properties of chemical elements depend on the charges of atomic nuclei.

table

Since the time of Mendeleev, the table he created has significantly changed and began to reflect almost all the functions and characteristics of the elements. The ability to use the table is essential for further study of chemistry. The modern table is presented in three forms:

• short - periods occupy two lines, and hydrogen is often referred to as group 7;
• long - isotopes and radioactive elements are taken out of the table;
• extra-long - each period occupies a separate line.

Rice. 3. Long modern table.

The short table is the most obsolete version that was canceled in 1989, but is still used in many textbooks. The long and extra long shapes are internationally recognized and used all over the world. Despite the established forms, scientists continue to improve the periodic system, offering the latest options.

What have we learned?

Mendeleev's periodic law and periodic system were formulated in 1871. Mendeleev identified the regularities of the properties of elements and ordered them on the basis of the relative atomic mass. As the mass increased, the properties of the elements changed and then repeated. Subsequently, the table was supplemented, and the law was adjusted in accordance with modern knowledge.

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Mendeleev's periodic law

D.I.Mendeleev's periodic law is a fundamental law that establishes periodic change properties of chemical elements depending on the increase in the charges of the nuclei of their atoms. I. Mendeleev in March 1869 when comparing the properties of all elements known at that time and the values ​​of their atomic masses. "The properties of simple bodies, as well as the shapes and properties of compounds of elements, and therefore the properties of the simple and complex bodies formed by them, are periodically dependent on their atomic weight." The graphic (tabular) expression of the periodic law is the periodic system of elements developed by Mendeleev.

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Figure 1. Dependence of the ionization energy of atoms on the ordinal number of the element

The energy of an atom's affinity for an electron, or simply its affinity for an electron, is the energy released during the attachment of an electron to a free atom E in its ground state with its transformation into a negative ion E− (the affinity of an atom for an electron is numerically equal, but opposite in sign of energy ionization of the corresponding isolated singly charged anion). The dependence of the electron affinity of an atom on the atomic number of an element is shown in Figure 2.

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Electronic configuration

Electronegativity is a fundamental chemical property of an atom, a quantitative characteristic of the ability of an atom in a molecule to attract common electron pairs to itself. The electronegativity of an atom depends on many factors, in particular on the valence state of the atom, the oxidation state, the coordination number, the nature of the ligands that make up the environment of the atom in the molecular system, and on some others. Figure 3 shows the dependence of electronegativity on the serial number of the element.

Figure 3. Pauling electronegativity scale

Recently, the so-called orbital electronegativity has been used more and more often to characterize electronegativity, which depends on the type of atomic orbital involved in the formation of a bond and on its electronic population, i.e. on whether the atomic orbital is occupied by a lone electron pair, is populated once by an unpaired electron, or is vacant. But, despite the well-known difficulties in the interpretation and definition of electronegativity, it always remains necessary for a qualitative description and prediction of the nature of bonds in a molecular system, including the binding energy, distribution of electronic charge, etc.

In periods, there is a general tendency towards an increase in electronegativity, and in subgroups, its decline. The smallest electronegativity is for s-elements of group I, the highest is for p-elements of group VII.

The periodicity in the change in the values ​​of the orbital atomic radii, depending on the ordinal number of the element, manifests itself quite clearly, and the main points here consist in the presence of very pronounced maxima corresponding to atoms of alkali metals, and the same minima corresponding to noble gases. A decrease in the values ​​of orbital atomic radii during the transition from an alkali metal to the corresponding (nearest) noble gas is, with the exception of the Li-Ne series, non-monotonic, especially when families of transition elements (metals) and lanthanides or actinides appear between the alkali metal and the noble gas. At large periods, a less sharp decrease in the radii is observed in the families of d and f elements, since the filling of the orbitals with electrons occurs in the pre-outer layer. In subgroups of elements, the radii of atoms and ions of the same type generally increase.

The oxidation state is an auxiliary conditional value for recording the processes of oxidation, reduction and redox reactions, the numerical value of the electric charge attributed to an atom in a molecule on the assumption that the electron pairs that make a connection are completely displaced towards more electronegative atoms.

Many elements are capable of exhibiting not one but several different oxidation states. For example, for chlorine, all oxidation states are known from −1 to +7, although even ones are very unstable, and for manganese, from +2 to +7. The highest values ​​of the oxidation state change periodically depending on the ordinal number of the element, but this periodicity is complex. In the simplest case, in the series of elements from an alkali metal to a noble gas, the highest oxidation state increases from +1 (RbF) to +8 (XeO4). In other cases, the highest oxidation state of the noble gas turns out to be lower (Kr + 4F4) than for the preceding halogen (Br + 7О4−). Therefore, on the curve of the periodic dependence of the highest oxidation state on the ordinal number of the element, the maxima fall either on the noble gas or on the halogen preceding it (the minima are always on the alkali metal). The exception is the Li-Ne series, in which high oxidation states are generally unknown neither for halogen (F), nor for noble gas (Ne), and the middle term of the series, nitrogen, has the highest value of the highest oxidation state; therefore, in the Li - Ne series, the change in the highest oxidation state turns out to be passing through a maximum.

In the general case, the increase in the highest oxidation state in the series of elements from an alkali metal to a halogen or to a noble gas is by no means monotonic, mainly due to the manifestation of high oxidation states with transition metals. For example, an increase in the highest oxidation state in the Rb-Xe series from +1 to +8 is "complicated" by the fact that for molybdenum, technetium and ruthenium, such high oxidation states are known as +6 (MoO3), +7 (Tc2O7), +8 (RuO4).

The change in the oxidative potentials of simple substances, depending on the ordinal number of the element, is also periodic. But it should be borne in mind that the oxidative potential of a simple substance is influenced by various factors, which sometimes need to be considered individually. Therefore, the periodicity in the change in oxidation potentials should be interpreted very carefully. You can find some definite sequences in the change in the oxidative potentials of simple substances. In particular, in a series of metals, when passing from alkaline to the elements following it, a decrease in oxidation potentials occurs. This is easily explained by an increase in the ionization energy of atoms with an increase in the number of removed valence electrons. Therefore, on the curve of the dependence of the oxidation potentials of simple substances on the ordinal number of the element, there are maxima corresponding to alkali metals.

DISCOVERY OF THE PERIODIC LAW

The periodic law was discovered by DI Mendeleev in the course of work on the text of the textbook "Fundamentals of Chemistry", when he encountered difficulties in systematizing factual material. By mid-February 1869, pondering the structure of the textbook, the scientist gradually came to the conclusion that the properties of simple substances and the atomic masses of elements are connected by a certain pattern.

The discovery of the periodic table of elements was not done by chance, it was the result of tremendous work, long and painstaking work that was expended both by Dmitry Ivanovich himself and by many chemists from among his predecessors and contemporaries. “When I began to finalize my classification of elements, I wrote on separate cards each element and its compounds, and then, arranging them in the order of groups and rows, I received the first visual table of the periodic law. But this was only the final chord, the result of all the previous work ... ”- said the scientist. Mendeleev emphasized that his discovery was the result that completed twenty years of thinking about the connections between elements, thinking from all sides of the relationship of elements.

On February 17 (March 1), the manuscript of the article, containing a table entitled "Experience of a system of elements based on their atomic weight and chemical similarity," was completed and sent to press with notes for typesetters and with the date "February 17, 1869". The announcement of Mendeleev's discovery was made by the editor of the Russian Chemical Society, Professor N.A. and Novgorod provinces.

In the first version of the system, the elements were arranged by the scientists in nineteen horizontal rows and six vertical columns. On February 17 (March 1), the opening of the periodic law was by no means completed, but just begun. Dmitry Ivanovich continued to develop and deepen it for almost three more years. In 1870, Mendeleev, in his Fundamentals of Chemistry, published the second version of the system (The Natural System of Elements): the horizontal columns of analogous elements turned into eight vertically arranged groups; the six vertical columns of the first variant turned into periods beginning with an alkali metal and ending with a halogen. Each period was divided into two rows; the elements of the different rows included in the group formed subgroups.

The essence of Mendeleev's discovery was that with an increase in the atomic mass of chemical elements, their properties change not monotonically, but periodically. After a certain number of elements of different properties, arranged in increasing atomic weight, the properties begin to repeat. The difference between the work of Mendeleev and the work of his predecessors was that Mendeleev had not one but two bases for the classification of elements - atomic mass and chemical similarity. In order for the periodicity to be fully observed, Mendeleev corrected the atomic masses of some elements, placed several elements in his system, contrary to the ideas accepted at that time about their similarity with others, left empty cells in the table where the elements that had not yet been discovered were to be located.

In 1871, on the basis of these works, Mendeleev formulated the Periodic Law, the form of which was somewhat improved over time.

The periodic table of elements had a great influence on the subsequent development of chemistry. It was not only the first natural classification of chemical elements, which showed that they form a harmonious system and are in close connection with each other, but also became a powerful tool for further research. At the time when Mendeleev compiled his table on the basis of the periodic law he discovered, many elements were still unknown. Mendeleev was not only convinced that there must be still unknown elements that would fill these places, but also predicted in advance the properties of such elements, based on their position among other elements of the periodic table. Over the next 15 years, Mendeleev's predictions were brilliantly confirmed; all three expected elements were discovered (Ga, Sc, Ge), which was the greatest triumph of the periodic law.

DI. Mendeleev submitted the manuscript "Experience of the system of elements based on their atomic weight and chemical similarity" // Presidential Library// A day in history http://www.prlib.ru/History/Pages/Item.aspx?itemid=1006

RUSSIAN CHEMICAL SOCIETY

The Russian Chemical Society is a scientific organization founded at St. Petersburg University in 1868 and was a voluntary association of Russian chemists.

The need to create the Society was announced at the 1st Congress of Russian Naturalists and Physicians, held in St. Petersburg at the end of December 1867 - early January 1868. At the Congress, the decision of the participants of the Chemical Section was announced:

“The Chemical Section has declared a unanimous desire to unite in the Chemical Society to communicate with the already established forces of Russian chemists. The section believes that this society will have members in all cities of Russia, and that its publication will include the works of all Russian chemists, printed in Russian. "

By this time, chemical societies had already been established in several European countries: the London Chemical Society (1841), the Chemical Society of France (1857), the German Chemical Society (1867); The American Chemical Society was founded in 1876.

The Charter of the Russian Chemical Society, drawn up mainly by D.I.Mendeleev, was approved by the Ministry of Public Education on October 26, 1868, and the first meeting of the Society took place on November 6, 1868. Initially, it included 35 chemists from St. Petersburg, Kazan, Moscow, Warsaw , Kiev, Kharkov and Odessa. N.N. Zinin became the first President of the RFB, and N.A.Menshutkin became the secretary. Members of the society paid membership fees (10 rubles per year), the admission of new members was carried out only on the recommendation of three existing ones. In the first year of its existence, the RCS grew from 35 to 60 members and continued to grow smoothly in subsequent years (129 - in 1879, 237 - in 1889, 293 - in 1899, 364 - in 1909, 565 - in 1917).

In 1869, the Russian Chemical Society got its own organ - the Journal of the Russian Chemical Society (ZhRHO); the magazine was published 9 times a year (monthly, except for the summer months). The editor of ZhRHO from 1869 to 1900 was N. A. Menshutkin, and from 1901 to 1930 - A. E. Favorsky.

In 1878, the Russian Chemical Society merged with the Russian Physical Society (founded in 1872) to form the Russian Physicochemical Society. The first Presidents of the RFCO were A.M.Butlerov (in 1878–1882) and D.I. Mendeleev (in 1883–1887). In connection with the merger since 1879 (from the 11th volume), the "Journal of the Russian Chemical Society" was renamed into the "Journal of the Russian Physicochemical Society". The frequency of publication was 10 issues per year; the journal consisted of two parts - chemical (ZhRHO) and physical (ZhRFO).

For the first time, many works of the classics of Russian chemistry were published on the pages of ZhRHO. We can especially note the work of D.I.Mendeleev on the creation and development of the periodic table of elements and A.M. Butlerov, associated with the development of his theory of structure organic compounds; research by N. A. Menshutkin, D. P. Konovalov, N. S. Kurnakov, L. A. Chugaev in the field of inorganic and physical chemistry; V. V. Markovnikova, E. E. Wagner, A. M. Zaitsev, S. N. Reformatsky, A. E. Favorsky, N. D. Zelinsky, S. V. Lebedev and A. E. Arbuzova in the field of organic chemistry. During the period from 1869 to 1930, 5067 original chemical studies were published in ZhRHO, abstracts and review articles on certain issues of chemistry, translations of the most interesting works from foreign journals were also published.

RFCO became the founder of the Mendeleev Congresses on General and Applied Chemistry; the first three congresses were held in St. Petersburg in 1907, 1911 and 1922. In 1919, the publication of ZhRFKhO was suspended and resumed only in 1924.