Periodic law, one of the fundamental laws of natural science, was discovered by the great Russian scientist D.I. Mendeleev in 1869. Initially, the law was formulated as follows: the properties of elements and their compounds are periodically dependent on their atomic weight(according to modern ideas - atomic mass).

The periodic law was presented as a classification of elements. On its basis, the elements were arranged into natural groups according to the totality of their properties. Special attention was paid to this point: guided by the properties of the elements, D.I. In a number of cases, Mendeleev even had to deviate from the sequential arrangement of elements in the Periodic Table strictly according to increasing atomic masses (atomic “weights”), for example, 18 Ar (39.9) and 19 K (39.1), 52 Te (127.6 ) and 53 1(126.9).

At the time of Mendeleev, the reason for the periodicity of the properties of elements was not known. However, the discoverer of the Periodic Law was confident that the reason should be sought in the structure of matter.

The discovery of the Periodic Law not only provided the foundation for chemical science, but also posed the task of elucidating the physical cause of periodicity. Chemical and the absolute majority physical properties elements are a periodic function of some independent, uniquely determined quantity inherent in each element and varying monotonically from element to element. Atomic mass (“atomic weight”) was accepted by Mendeleev as such a value.

Only when, thanks to the successes of physics, much more was known about the structure of the atom than at the time of the discovery and establishment of the periodic law, its true meaning and the reasons for periodicity became clear. From element to element according to the Periodic Table, the charge of the nucleus of an element’s atom changes, which is determined by the number of protons. In the Periodic Table, this number coincides with the atomic number of the element. Since the atom is electrically neutral, the charge of the nucleus (in units of electron charge) is equal to the number of electrons in the electron shell of the atom. Increase serial number element per unit means that one proton has been added to the nucleus of the atom, and accordingly one electron has been added to the electron shell. Since the properties of elements, especially chemical ones, are determined mainly by the electrons of the outer quantum layer, the reason for the periodicity of the properties is the periodic nature of the filling of the space around the nucleus with electrons. The factor that determines the structure of the electronic shells of atoms, and therefore the properties of elements, is the charge of the atomic nucleus. Therefore, the modern formulation of the periodic law is as follows: The properties of elements and their compounds are periodically dependent on the charge of the nucleus of the element’s atom.

The atomic mass of an element is determined by the total number of nucleons (protons and neutrons) in the isotope nuclei of this element and the isotopic composition of the element. The change in atomic mass is mainly proportional to the charge of the nucleus. Therefore, Mendeleev's formulation of the Periodic Law, with few exceptions, correctly reflects the arrangement of elements in the Periodic Table, but does not reveal the reason for periodicity.

According to the Pauli principle, the number of possible electronic states in quantum levels and sublevels is limited by the number of combinations of non-repeating sets of four quantum numbers P, /, T And s, and this determines the capacity of quantum levels and sublevels (see Table 2.1). If the atom is not excited, electrons fill those orbitals whose energy is minimal.

The periodic table would be simpler if the energy in multielectron atoms, as in the hydrogen atom, was determined by the principal quantum number. Then, in accordance with the capacity of the quantum layers, the periods would consist of 2, 8, 18, 32, 50, etc. elements, and noble gases with a complete quantum level would have numbers 2, 10, 28, 60, 110... However, due to electron-electron interaction, this sequence is disrupted. From the IV period, the filling of a new quantum layer, which in the Periodic System corresponds to the beginning of a new period, begins with the incomplete pre-external III quantum level, and from the VI period - with incomplete IV and V quantum levels, etc. Therefore, noble gases - elements after which the construction of a new quantum level (and a new period) begins - contain only 8 electrons on the outer quantum layer and have numbers 2, 10, 18, 36, 54, and 86. Accordingly, the periods cover 2. 8, 8, 18, 18 and 32 elements.

The periodic law does not have a specific mathematical expression. It is represented in the form of a periodic table. There are several versions of such a table, but all of them are presented in one form or another as structure diagrams of the atomic structure of any element. It becomes possible to establish the electronic structure of any atom not only on the basis of the known sequence of filling sublevels or the Klechkovsky rule, but also on the basis of the table itself: the position of an element in the table uniquely reflects the electronic structure of its atoms. The distribution of elements by periods and subgroups exactly corresponds to the distribution of electrons of the atoms of these elements by levels and sublevels of the electron shell.

DI. Mendeleev formulated the Periodic Law in 1869, which was based on one of the main characteristics atom - atomic mass. The subsequent development of the Periodic Law, namely, the acquisition of a large amount of experimental data, somewhat changed the original formulation of the law, but these changes do not contradict the main meaning laid down by D.I. Mendeleev. These changes only gave the law and the Periodic Table scientific validity and confirmation of correctness.

Modern formulation of the Periodic Law by D.I. Mendeleev is as follows: the properties of chemical elements, as well as the properties and forms of compounds of elements, are periodically dependent on the magnitude of the charge of the nuclei of their atoms.

Structure of the Periodic Table of Chemical Elements D.I. Mendeleev

By present opinion it is known a large number of interpretations of the Periodic Table, but the most popular is with short (small) and long (large) periods. Horizontal rows are called periods (they contain elements with sequential filling of the same energy level), and vertical columns are called groups (they contain elements that have the same number of valence electrons - chemical analogues). Also, all elements can be divided into blocks according to the type of external (valence) orbital: s-, p-, d-, f-elements.

There are a total of 7 periods in the system (table), and the number of the period (indicated by an Arabic numeral) is equal to the number of electronic layers in the atom of the element, the number of the external (valence) energy level, and the value of the principal quantum number for the highest energy level. Each period (except the first) begins with an s-element - an active alkali metal and ends with an inert gas, preceded by a p-element - an active non-metal (halogen). If you move through the period from left to right, then with an increase in the charge of the nuclei of atoms of chemical elements of small periods, the number of electrons at the external energy level will increase, as a result of which the properties of the elements change - from typically metallic (since at the beginning of the period there is an active alkali metal), through amphoteric (the element exhibits the properties of both metals and non-metals) to non-metallic (the active non-metal is halogen at the end of the period), i.e. metallic properties gradually weaken and non-metallic properties increase.

In large periods, as the charge of nuclei increases, the filling of electrons is more difficult, which explains a more complex change in the properties of elements compared to elements of small periods. Thus, in even rows of long periods, as the charge of the nucleus increases, the number of electrons in the outer energy level remains constant and equal to 2 or 1. Therefore, while the level next to the outer (second from the outside) is filled with electrons, the properties of the elements in the even rows change slowly. When moving to odd series, with increasing nuclear charge, the number of electrons in the external energy level increases (from 1 to 8), the properties of the elements change in the same way as in small periods.

Vertical columns in the Periodic Table are groups of elements with similar electronic structure and being chemical analogues. Groups are designated by Roman numerals from I to VIII. There are main (A) and secondary (B) subgroups, the first of which contain s- and p-elements, the second - d-elements.

The number A of the subgroup shows the number of electrons in the outer energy level (the number of valence electrons). For B-subgroup elements, there is no direct connection between the group number and the number of electrons in the outer energy level. In A-subgroups, the metallic properties of elements increase, and non-metallic properties decrease with increasing charge of the nucleus of the element’s atom.

There is a relationship between the position of elements in the Periodic Table and the structure of their atoms:

- atoms of all elements of the same period have an equal number of energy levels, partially or completely filled with electrons;

- atoms of all elements of the A subgroups have an equal number of electrons at the outer energy level.

Periodic properties of elements

The similarity of the physicochemical and chemical properties of atoms is due to the similarity of their electronic configurations, and main role plays the distribution of electrons in the outer atomic orbital. This manifests itself in the periodic appearance, as the charge of the atomic nucleus increases, of elements with similar properties. Such properties are called periodic, among which the most important are:

1. Number of electrons in the outer electron shell ( population – w). In short periods with increasing nuclear charge w the outer electron shell monotonically increases from 1 to 2 (1st period), from 1 to 8 (2nd and 3rd periods). In large periods during the first 12 elements w does not exceed 2, and then up to 8.

2. Atomic and ionic radii(r), defined as the average radii of an atom or ion, found from experimental data on interatomic distances in different compounds. According to the period, the atomic radius decreases (gradually adding electrons are described by orbitals with almost equal characteristics; according to the group, the atomic radius increases as the number of electron layers increases (Fig. 1.).

Rice. 1. Periodic change in atomic radius

The same patterns are observed for the ionic radius. It should be noted that the ionic radius of the cation (positively charged ion) is greater than the atomic radius, which in turn is greater than the ionic radius of the anion (negatively charged ion).

3. Ionization energy(E and) is the amount of energy required to remove an electron from an atom, i.e. the energy required to transform a neutral atom into a positively charged ion (cation).

E 0 - → E + + E and

E and is measured in electronvolts (eV) per atom. Within the group of the Periodic Table, the values ​​of ionization energy of atoms decrease with increasing charges of the atomic nuclei of elements. All electrons can be sequentially removed from atoms of chemical elements, giving discrete values E and. Moreover, E and 1< Е и 2 < Е и 3 <….Энергии ионизации отражают дискретность структуры электронных слоев и оболочек атомов химических элементов.

4. Electron affinity(E e) – the amount of energy released when an additional electron is added to an atom, i.e. process energy

E 0 + → E —

E e is also expressed in eV and, like E, it depends on the radius of the atom, therefore the nature of the change in E e across periods and groups of the Periodic System is close to the nature of the change in the atomic radius. Group VII p-elements have the highest electron affinity.

5. Regenerative activity(VA) – the ability of an atom to give an electron to another atom. Quantitative measure – E and. If E increases, then BA decreases and vice versa.

6. Oxidative activity(OA) – the ability of an atom to attach an electron from another atom. Quantitative measure E e. If E e increases, then OA also increases and vice versa.

7. Shielding effect– reducing the impact of the positive charge of the nucleus on a given electron due to the presence of other electrons between it and the nucleus. Shielding increases with the number of electron layers in an atom and reduces the attraction of outer electrons to the nucleus. The opposite of shielding penetration effect, due to the fact that the electron can be located at any point in atomic space. The penetration effect increases the strength of the bond between the electron and the nucleus.

8. Oxidation state (oxidation number)– the imaginary charge of an atom of an element in a compound, which is determined from the assumption of the ionic structure of the substance. The group number of the Periodic Table indicates the highest positive oxidation state that elements of a given group can have in their compounds. Exceptions are metals of the copper subgroup, oxygen, fluorine, bromine, metals of the iron family and other elements of group VIII. As the nuclear charge increases in a period, the maximum positive oxidation state increases.

9. Electronegativity, compositions of higher hydrogen and oxygen compounds, thermodynamic, electrolytic properties, etc.

Examples of problem solving

EXAMPLE 1

Exercise	Characterize the element (Z=23) and the properties of its compounds (oxides and hydroxides) using the electronic formula: family, period, group, number of valence electrons, electron graphic formula for valence electrons in the ground and excited states, main oxidation states (maximum and minimum ), formulas of oxides and hydroxides.
Solution	23 V 1s 2 2s 2 2p 6 3s 3 3p 6 3d 3 4s 2 d-element, metal, is in the ;-th period, in the V group, in the subgroup. Valence electrons 3d 3 4s 2. Oxides VO, V 2 O 3, VO 2, V 2 O 5. Hydroxides V(OH)2, V(OH)3, VO(OH)2, HVO3. Ground state Excited state The minimum oxidation state is “+2”, the maximum is “+5”.

By the time the periodic law was discovered, 63 chemical elements were known and the properties of their various compounds were described.

Works of predecessors D.I. Mendeleev:

1. Berzelius classification, which has not lost its relevance today (metals, non-metals)

2. Döbereiner triads (eg lithium, sodium, potassium)

4. Shankurtur spiral-axis

5. Meyer curve

Participation of D.I. Mendeleev at the International Chemical Congress in Karslruhe (1860), where the ideas of atomism and the concept of “Atomic” weight, which is now known as “relative atomic mass,” were established.

Personal qualities of the great Russian scientist D.I. Mendeleev.

The brilliant Russian chemist was distinguished by his encyclopedic knowledge, scrupulous chemical experiment, the greatest scientific intuition, confidence in the truth of his position and hence the undaunted risk in defending this truth. DI. Mendeleev was a great and wonderful citizen of the Russian land.

D.I. Mendeleev arranged all the chemical elements known to him in a long chain in increasing order of their atomic weights and noted segments in it - periods in which the properties of the elements and the substances formed by them changed in a similar way, namely:

1). Metallic properties weakened;

2) Non-metallic properties were enhanced;

3) The degree of oxidation in higher oxides increased from +1 to +7(+8);

4).The degree of oxidation of elements in hydroxides, solid salt-like compounds of metals with hydrogen increased from +1 to +3, and then in volatile hydrogen compounds from -4 to -1;

5) Oxides from basic through amphoteric were replaced by acidic ones;

6) Hydroxides from alkalis, through amphoteric ones, were replaced by acids.

The conclusion of his work was the first formulation of the periodic law (March 1, 1869): the properties of chemical elements and the substances formed by them are periodically dependent on their relative atomic masses.

Periodic law and atomic structure.

The formulation of the periodic law given by Mendeleev was inaccurate and incomplete, because it reflected the state of science at a time when the complex structure of the atom was not yet known. Therefore, the modern formulation of the periodic law sounds different: the properties of chemical elements and the substances formed by them are periodically dependent on the charge of their atomic nuclei.

Periodic table and atomic structure.

The periodic table is a graphical representation of the periodic law.

Each designation in the periodic table reflects some feature or pattern in the structure of the atoms of the elements:

Physical meaning of element number, period, group;

Reasons for changes in the properties of elements and substances formed by them horizontally (in periods) and vertically (in groups).

Within the same period, metallic properties weaken, and non-metallic properties increase, because:

1) The charges of atomic nuclei increase;

2) The number of electrons at the external level increases;

3) The number of energy levels is constant;

4) The radius of the atom decreases

Within the same group (in the main subgroup), metallic properties increase, non-metallic properties weaken, because:

1). The charges of atomic nuclei increase;

2). The number of electrons in the outer level is constant;

3). The number of energy levels increases;

4). The radius of the atom increases

As a result of this, a cause-and-effect formulation of the periodic law was given: the properties of chemical elements and the substances formed by them are periodically dependent on changes in the external electronic structures of their atoms.

The meaning of the periodic law and the periodic system:

1. Allowed us to establish the relationship between elements and combine them by properties;

2. Arrange chemical elements in natural sequence;

3. Reveal the periodicity, i.e. repeatability of the general properties of individual elements and their compounds;

4. Correct and clarify the relative atomic masses of individual elements (for beryllium from 13 to 9);

5. Correct and clarify the oxidation states of individual elements (beryllium +3 to +2)

6. Predict and describe the properties, indicate the path to the discovery of yet undiscovered elements (scandium, gallium, germanium)

Using the table, we compare the two leading theories of chemistry.

Philosophical foundations of community	Periodic law of D.I.Mendeleev	Theory of organic compounds A.M. Butlerov
1. 1. Opening time	1869	1861
II. Prerequisites. 1. Accumulation of factual material 2. 2. Work of predecessors 3. Congress of chemists in Karlsruhe (1860) 4. Personal qualities.	By the time the periodic law was discovered, 63 chemical elements were known and the properties of their numerous compounds were described.	Many tens and hundreds of thousands of organic compounds are known, consisting of only a few elements: carbon, hydrogen, oxygen, and, less commonly, nitrogen, phosphorus and sulfur.
- J. Berzelius (metals and non-metals) - I.V. Debereiner (triads) - D.A.R. Newlands (octaves) - L. Meyer	- J. Bercellius, J. Liebig, J. Dumas (theory of radicals); -J. Dumas, C. Gerard, O. Laurent (theory of types); - J. Berzelius introduced the term “isomerism” into practice; -F.Vehler, N.N. Zinin, M. Berthelot, A. Butlerov himself (synthesis of organic substances, collapse of vitalism); -F.A. Kukule (structure of benzene)
DI. Mendeleev was present as an observer	A.M. Butlerov did not participate, but actively studied the materials of the congress. However, he took part in the congress of doctors and natural scientists in Speyer (1861), where he made a report “On the structure of organic bodies”
Both authors were distinguished from other chemists by: the encyclopedic nature of chemical knowledge, the ability to analyze and generalize facts, scientific forecasting, Russian mentality and Russian patriotism.
III. The role of practice in the development of theory	DI. Mendeleev predicts and indicates the path to the discovery of gallium, scandium and germanium, still unknown to science	A.M. Butlerov predicts and explains the isomerism of many organic compounds. He carries out many syntheses himself.

Test on the topic

Periodic law and periodic system of elements D.I. Mendeleev

1. How do the radii of atoms change in a period:

2. How do the radii of atoms change in the main subgroups:

a) increase b) decrease c) do not change

3. How to determine the number of energy levels in an atom of an element:

a) by element serial number b) by group number

c) by row number d) by period number

4. How is the place of a chemical element in the periodic table determined by D.I. Mendeleev:

a) the number of electrons in the outer level b) the number of neutrons in the nucleus

c) charge of the atomic nucleus d) atomic mass

5. How many energy levels does the scandium atom have: a) 1 b) 2 c) 3 d) 4

6. What determines the properties of chemical elements:

a) the value of the relative atomic mass b) the number of electrons in the outer layer

c) charge of the atomic nucleus d) number of valence electrons

7. How do the chemical properties of elements change during the period:

a) metallic ones are amplified b) non-metallic ones are amplified

c) do not change d) non-metallic weaken

8. Indicate the element that heads the large period of the periodic table of elements: a) Cu (No. 29) b) Ag (No. 47) c) Rb (No. 37) d) Au (No. 79)

9. Which element has the most pronounced metallic properties:

a) Magnesium b) Aluminum c) Silicon

10. Which element has the most pronounced non-metallic properties:

a) Oxygen b) Sulfur c) Selenium

11.What is the main reason for changes in the properties of elements in periods:

a) in an increase in atomic masses

b) in a gradual increase in the number of electrons at the external energy level

c) in increasing the number of electrons in an atom

d) in increasing the number of neutrons in the nucleus

12. Which element heads the main subgroup of the fifth group:

a) vanadium b) nitrogen c) phosphorus d) arsenic

13.What is the number of orbitals at the d-sublevel: a)1 b)3 c)7 d)5

14. How do atoms of isotopes of the same element differ:

a) the number of protons b) the number of neutrons c) the number of electrons d) the charge of the nucleus

15. What is an orbital:

a) a certain energy level at which the electron is located

b) the space around the nucleus where the electron is located

c) the space around the nucleus, where the probability of finding an electron is greatest

d) the trajectory along which the electron moves

16. In which orbital does the electron have the highest energy: a) 1s b) 2s c) 3s d) 2p

17. Determine what element 1s 2 2s 2 2p 1 is: a) No. 1 b) No. 3 c) No. 5 d) No. 7

18. What is the number of neutrons in an atom +15 31 R a)31 b)16 c)15 d)46

19. Which element has the structure of the outer electron layer ...3s 2 p 6:

a) neon b) chlorine c) argon d) sulfur

20. Based on the electronic formula, determine what properties the element 1s 2 2s 2 2p 5 has:

a) metal b) non-metal c) amphoteric element d) inert element

21. How many chemical elements are in the sixth period: a)8 b)18 c)30 d)32

22. What is the mass number of nitrogen +7 N which contains 8 neutrons:

a)14 b)15 c)16 d)17

23. An element whose atomic nucleus contains 26 protons: a) S b) Cu c) Fe d) Ca

From your first chemistry lessons you used D.I. Mendeleev’s table. It clearly demonstrates that all the chemical elements that form the substances of the world around us are interconnected and obey general laws, that is, they represent a single whole - a system of chemical elements. Therefore, in modern science, D.I. Mendeleev’s table is called the Periodic Table of Chemical Elements.

Why “periodic” is also clear to you, since the general patterns in changes in the properties of atoms, simple and complex substances formed by chemical elements are repeated in this system at certain intervals - periods. Some of these patterns shown in Table 1 are already known to you.

Thus, all chemical elements existing in the world are subject to a single, objectively valid Periodic Law in nature, the graphic representation of which is the Periodic Table of Elements. This law and system are named after the great Russian chemist D.I. Mendeleev.

D.I. Mendeleev came to the discovery of the Periodic Law by comparing the properties and relative atomic masses of chemical elements. To do this, D.I. Mendeleev wrote down on a card for each chemical element: the symbol of the element, the value of the relative atomic mass (at the time of D.I. Mendeleev this value was called atomic weight), the formulas and nature of the higher oxide and hydroxide. He arranged 63 chemical elements known by that time into one chain in increasing order of their relative atomic masses (Fig. 1) and analyzed this set of elements, trying to find certain patterns in it. As a result of intense creative work, he discovered that there are intervals in this chain - periods in which the properties of the elements and the substances formed by them change in a similar way (Fig. 2).

Rice. 1.
Cards of elements arranged in increasing order of their relative atomic masses

Rice. 2.
Cards of elements arranged in order of periodic changes in the properties of elements and substances formed by them

Laboratory experiment No. 2
Modeling the construction of the Periodic Table of D. I. Mendeleev

Model the construction of the Periodic Table of D.I. Mendeleev. To do this, prepare 20 cards measuring 6 x 10 cm for elements with serial numbers from 1st to 20th. On each card, indicate the following information about the element: chemical symbol, name, relative atomic mass, formula of higher oxide, hydroxide (indicate their nature in parentheses - basic, acidic or amphoteric), formula of volatile hydrogen compound (for non-metals). Shuffle the cards and then arrange them in a row in order of increasing relative atomic masses of the elements. Place similar elements from 1st to 18th under each other: hydrogen above lithium and potassium under sodium, respectively, calcium under magnesium, helium under neon. Formulate the pattern you have identified in the form of a law. Note the discrepancy between the relative atomic masses of argon and potassium and their location in terms of the common properties of the elements. Explain the reason for this phenomenon.

Let us list once again, using modern terms, the regular changes in properties that manifest themselves within periods:

• metallic properties weaken;
• non-metallic properties are enhanced;
• the degree of oxidation of elements in higher oxides increases from +1 to +8;
• the oxidation degree of elements in volatile hydrogen compounds increases from -4 to -1;
• oxides from basic through amphoteric are replaced by acidic ones;
• hydroxides from alkalis through amphoteric hydroxides are replaced by oxygen-containing acids.

Based on these observations, D.I. Mendeleev made a conclusion in 1869 - he formulated the Periodic Law, which, using modern terms, sounds like this:

Systematizing chemical elements based on their relative atomic masses, D. I. Mendeleev also paid great attention to the properties of the elements and the substances formed by them, distributing elements with similar properties into vertical columns - groups. Sometimes, in violation of the pattern he had identified, he placed heavier elements in front of elements with lower relative atomic masses. For example, he wrote cobalt in his table before nickel, tellurium before iodine, and when inert (noble) gases were discovered, argon before potassium. D.I. Mendeleev considered this order of arrangement necessary because otherwise these elements would fall into groups of elements dissimilar to them in properties. So, in particular, the alkali metal potassium would fall into the group of inert gases, and the inert gas argon would fall into the group of alkali metals.

D.I. Mendeleev could not explain these exceptions to the general rule, as well as the reason for the periodicity in changes in the properties of elements and the substances formed by them. However, he foresaw that this reason lay in the complex structure of the atom. It was the scientific intuition of D.I. Mendeleev that allowed him to construct a system of chemical elements not in the order of increasing their relative atomic masses, but in the order of increasing charges of their atomic nuclei. The fact that the properties of elements are determined precisely by the charges of their atomic nuclei is eloquently demonstrated by the existence of isotopes that you met last year (remember what they are, give examples of isotopes known to you).

In accordance with modern ideas about the structure of the atom, the basis for the classification of chemical elements is the charges of their atomic nuclei, and the modern formulation of the Periodic Law is as follows:

The periodicity in changes in the properties of elements and their compounds is explained by the periodic repetition in the structure of the external energy levels of their atoms. It is the number of energy levels, the total number of electrons located on them and the number of electrons at the outer level that reflect the symbolism adopted in the Periodic System, that is, they reveal the physical meaning of the element’s serial number, period number and group number (what does it consist of?).

The structure of the atom makes it possible to explain the reasons for changes in the metallic and non-metallic properties of elements in periods and groups.

Consequently, the Periodic Law and the Periodic System of D.I. Mendeleev summarize information about chemical elements and the substances formed by them and explain the periodicity in changes in their properties and the reason for the similarity of the properties of elements of the same group.

These two most important meanings of the Periodic Law and the Periodic System of D.I. Mendeleev are complemented by one more, which is the ability to predict, i.e. predict, describe properties and indicate ways of discovering new chemical elements. Already at the stage of creating the Periodic Table, D.I. Mendeleev made a number of predictions about the properties of elements not yet known at that time and indicated the ways of their discovery. In the table he created, D.I. Mendeleev left empty cells for these elements (Fig. 3).

Rice. 3.
Periodic table of elements proposed by D. I. Mendeleev

Vivid examples of the predictive power of the Periodic Law were the subsequent discoveries of elements: in 1875, the Frenchman Lecoq de Boisbaudran discovered gallium, predicted by D. I. Mendeleev five years earlier as an element called “ekaaluminum” (eka - next); in 1879, the Swede L. Nilsson discovered the “ekabor” according to D. I. Mendeleev; in 1886 by the German K. Winkler - “exasilicon” according to D. I. Mendeleev (determine the modern names of these elements from D. I. Mendeleev’s table). How accurate D.I. Mendeleev was in his predictions is illustrated by the data in Table 2.

table 2
Predicted and experimentally discovered properties of germanium

Predicted by D.I. Mendeleev in 1871	Established by K. Winkler in 1886
Relative atomic mass is close to 72	Relative atomic mass 72.6
Gray refractory metal	Gray refractory metal
Metal density is about 5.5 g/cm 3	Metal density 5.35 g/cm 3
Oxide formula E0 2	Ge0 2 oxide formula
Oxide density is about 4.7 g/cm3	Oxide density 4.7 g/cm3
The oxide will be reduced to metal quite easily	Ge0 2 oxide is reduced to metal when heated in a hydrogen stream
Chloride ES1 4 should be a liquid with a boiling point of about 90 °C and a density of about 1.9 g/cm3	Germanium (IV) chloride GeCl 4 is a liquid with a boiling point of 83 ° C and a density of 1.887 g/cm 3

Scientists who discovered new elements highly appreciated the discovery of the Russian scientist: “There can hardly be a more striking proof of the validity of the doctrine of the periodicity of elements than the discovery of the still hypothetical eca-silicon; it constitutes, of course, more than a simple confirmation of a bold theory - it marks an outstanding expansion of the chemical field of vision, a giant step in the field of knowledge” (K. Winkler).

The American scientists who discovered element No. 101 gave it the name “mendelevium” in recognition of the great Russian chemist Dmitri Mendeleev, who was the first to use the Periodic Table of Elements to predict the properties of then undiscovered elements.

You met in 8th grade and will be using a form of the periodic table this year called the short period form. However, in specialized classes and in higher education, a different form is predominantly used - the long-period version. Compare them. What are the same and what are different about these two forms of the Periodic Table?

New words and concepts

1. Periodic law of D. I. Mendeleev.
2. The periodic table of chemical elements by D.I. Mendeleev is a graphical representation of the Periodic Law.
3. The physical meaning of element number, period number and group number.
4. Patterns of changes in the properties of elements in periods and groups.
5. The meaning of the Periodic Law and the Periodic Table of Chemical Elements by D. I. Mendeleev.

Tasks for independent work

1. Prove that the Periodic Law of D.I. Mendeleev, like any other law of nature, performs explanatory, generalizing and predictive functions. Give examples illustrating these functions of other laws known to you from chemistry, physics and biology courses.
2. Name a chemical element in the atom of which electrons are arranged in levels according to a series of numbers: 2, 5. What simple substance does this element form? What is the formula of its hydrogen compound and what is it called? What is the formula of the highest oxide of this element, what is its character? Write down the reaction equations characterizing the properties of this oxide.
3. Beryllium was previously classified as a Group III element, and its relative atomic mass was considered to be 13.5. Why did D.I. Mendeleev move it to group II and correct the atomic mass of beryllium from 13.5 to 9?
4. Write the reaction equations between a simple substance formed by a chemical element, in an atom of which electrons are distributed among energy levels according to a series of numbers: 2, 8, 8, 2, and simple substances formed by elements No. 7 and No. 8 in the Periodic Table. What type of chemical bond is present in the reaction products? What crystal structure do the original simple substances and the products of their interaction have?
5. Arrange the following elements in order of increasing metallic properties: As, Sb, N, P, Bi. Justify the resulting series based on the structure of the atoms of these elements.
6. Arrange the following elements in order of increasing non-metallic properties: Si, Al, P, S, Cl, Mg, Na. Justify the resulting series based on the structure of the atoms of these elements.
7. Arrange in order of weakening acidic properties the oxides whose formulas are: SiO 2, P 2 O 5, Al 2 O 3, Na 2 O, MgO, Cl 2 O 7. Justify the resulting series. Write down the formulas of the hydroxides corresponding to these oxides. How does their acidic character change in the series you proposed?
8. Write the formulas of boron, beryllium and lithium oxides and arrange them in ascending order of their main properties. Write down the formulas of the hydroxides corresponding to these oxides. What is their chemical nature?
9. What are isotopes? How did the discovery of isotopes contribute to the development of the Periodic Law?
10. Why do the charges of the atomic nuclei of elements in the Periodic Table of D.I. Mendeleev change monotonically, that is, the charge of the nucleus of each subsequent element increases by one compared to the charge of the atomic nucleus of the previous element, and the properties of the elements and the substances they form change periodically?
11. Give three formulations of the Periodic Law, in which the relative atomic mass, charge of the atomic nucleus and the structure of external energy levels in the electron shell of the atom are taken as the basis for the systematization of chemical elements.

Periodic law of chemical elements- a fundamental law of nature, reflecting the periodic change in the properties of chemical elements as the charges of the nuclei of their atoms increase. Opened on March 1 (February 17, Old Style) 1869 D.I. Mendeleev. On this day, he compiled a table called “Experience of a system of elements based on their atomic weight and chemical similarity.” The final formulation of the periodic law was given by Mendeleev in July 1871. It read:

“The properties of the elements, and therefore the properties of the simple and complex bodies they form, are periodically dependent on their atomic weight.”

Mendeleev's formulation of the periodic law existed in science for a little over 40 years. It was revised due to outstanding achievements in physics, mainly the development of the nuclear model of the atom (see Atom). It turned out that the charge of the nucleus of an atom (Z) is numerically equal to the serial number of the corresponding element in the periodic table, and the filling of the electronic shells and subshells of atoms, depending on Z, occurs in such a way that similar electronic configurations of atoms are periodically repeated (see Periodic system of chemical elements). Therefore, the modern formulation of the periodic law is as follows: the properties of elements, simple substances and their compounds are periodically dependent on the charges of atomic nuclei.
Unlike other fundamental laws of nature, such as the law of universal gravitation or the law of equivalence of mass and energy, the periodic law cannot be written in the form of any general equation or formula. Its visual reflection is the periodic table of elements. However, Mendeleev himself and other scientists made attempts to find mathematical equation of the periodic law of chemical elements. These attempts were crowned with success only after the development of the theory of atomic structure. But they concern only the establishment of the quantitative dependence of the order of distribution of electrons in shells and subshells on the charges of atomic nuclei.
Thus, by solving the Schrödinger equation, one can calculate how electrons are distributed in atoms with different Z values. And therefore, the basic equation of quantum mechanics is, as it were, one of the quantitative expressions of the periodic law.
Or, for example, another equation: Z„, = „+,Z - - (21 + 1)2 - >n,(2t + 1) +
1
+ m„where „+,Z = - (n + 1+ 1)" +
+(+1+ 1. 2k(p+O 1
2 2 6
Despite its bulkiness, it is not that difficult. The letters u, 1, t, and m are nothing more than the main, orbital, magnetic and spin quantum numbers (see Atom). The equation allows us to calculate at what value of Z (the atomic number of an element) an electron appears in an atom, the state of which is described by a given combination of four quantum numbers. Substituting possible combinations of u, 1, m, and m into this equation, we get a set of different values ​​of Z. If these values ​​are arranged in the sequence of the natural numbers 1, 2, 3, 4, 5, ..., then, in their In turn, a clear scheme is obtained for constructing the electronic configurations of atoms as Z increases. Thus, this equation is also a kind of quantitative expression of the periodic law. Try to solve this equation yourself for all elements of the periodic table (you will learn how the values ​​of u, 1; m, and m are related to each other in the article Atom).

The periodic law is a universal law for the entire Universe. It has power wherever atoms exist. But not only the electronic structures of atoms change periodically. The structure and properties of atomic nuclei also obey a peculiar periodic law. In nuclei consisting of neutrons and protons, there are neutron and proton shells, the filling of which is periodic. There are even known attempts to construct a periodic system of atomic nuclei.