The first law of Mendeleev. Periodic law, periodic system of chemical elements of Mendeleev and the structure of the atom

06.08.202027.05.2017

First option Periodic table of elements was published by Dmitri Ivanovich Mendeleev in 1869 and was called "The Experience of a System of Elements".

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

The properties of simple bodies, as well as the forms and properties of compounds of elements, are in a periodic dependence on the magnitude of the atomic weights of the elements.

The current wording of the law reads as follows:

The properties of chemical elements, simple substances, as well as the composition and properties of compounds are in a periodic dependence on the values of the charges of the nuclei of atoms.

Graphic image periodic law is the periodic table.

The cell of each element indicates its most important characteristics.

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 system 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 Chalcogens.

VII-a Halogens.

VIII-a Noble (inert) gases.

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

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

Parameter	Down the group	By period to the right
Core charge	is increasing	is increasing
Number of valence electrons	Does not change	is increasing
Number of energy levels	is increasing	Does not change
Atom radius	is increasing	Decreases
Electronegativity	Decreases	is increasing
Metal properties	Are increasing	Decrease
Oxidation state in higher oxide	Does not change	is increasing
The degree of oxidation in hydrogen compounds (for elements of groups IV-VII)	Does not change	is increasing

Modern periodic table of chemical elements of Mendeleev.

Here the reader will find information about one of the most important laws ever discovered by man in the scientific field - the periodic law of Mendeleev Dmitry Ivanovich. You will get acquainted with its meaning and influence on chemistry, will be considered general provisions, characteristics and details of the periodic law, the history of the discovery and the main provisions.

What is the periodic law

The periodic law is natural law of a fundamental nature, which was first discovered by D. I. Mendeleev back in 1869, and the discovery itself occurred due to a comparison of the properties of some chemical elements and the atomic mass values known at that time.

Mendeleev argued that, according to his law, simple and complex bodies and the various combinations of elements depend on their dependence of the periodic type and on the weight of their atom.

The periodic law is unique in its kind and this is due to the fact that it is not expressed by mathematical equations, unlike other fundamental laws of nature and the universe. Graphically, it finds its expression in the periodic table of chemical elements.

Discovery history

The discovery of the periodic law took place in 1869, but attempts to systematize all known x elements began long before that.

The first attempt to create such a system was made by I. V. Debereiner in 1829. He classified all the chemical elements known to him into triads, interconnected by the proximity of half the sum of the atomic masses included in this group of three components. Following Debereiner, an attempt was made to create a unique table of classification of the elements by A. de Chancourtua, he called his system the "earth spiral", and after him the Newlands octave was compiled by John Newlands. In 1864, almost simultaneously, William Olding and Lothar Meyer published independently created tables.

The periodic law was presented to the scientific community for review on March 8, 1869, and this happened during a meeting of the Russian X-th Society. Mendeleev Dmitry Ivanovich announced his discovery in front of everyone and in the same year Mendeleev's textbook "Fundamentals of Chemistry" was published, where the periodic table created by him was shown for the first time. A year later, in 1870, he wrote an article and submitted it for review to the RCS, where the concept of the periodic law was first used. In 1871, Mendeleev gave an exhaustive description of his research in his famous article on the periodic validity of chemical elements.

An invaluable contribution to the development of chemistry

The value of the periodic law is incredibly great for the scientific community around the world. This is due to the fact that its discovery gave a powerful impetus to the development of both chemistry and other natural sciences, such as physics and biology. The relationship of elements with their qualitative chemical and physical characteristics was open, and this also made it possible to understand the essence of the construction of all elements according to one principle and gave rise to the modern formulation of the concepts of chemical elements, to concretize knowledge about substances of complex and simple structure.

The use of the periodic law made it possible to solve the problem of chemical prediction, to determine the cause of the behavior of known chemical elements. Atomic physics, including nuclear energy, became possible as a result of the same law. In turn, these sciences made it possible to expand the horizons of the essence of this law and delve into its understanding.

Chemical properties of the elements of the periodic system

In fact, the chemical elements are interconnected by the characteristics inherent in them in the state of both a free atom and an ion, solvated or hydrated, in a simple substance and in the form that their numerous compounds can form. However, x-th properties usually consist in two phenomena: properties characteristic of an atom in a free state, and a simple substance. This kind of properties includes many of their types, but the most important are:

Atomic ionization and its energy, depending on the position of the element in the table, its ordinal number.
The energy relationship of the atom and electron, which, like atomic ionization, depends on the location of the element in the periodic table.
The electronegativity of an atom, which does not have a constant value, but can change depending on various factors.
The radii of atoms and ions - here, as a rule, empirical data are used, which is associated with the wave nature of electrons in a state of motion.
Atomization of simple substances - a description of the ability of an element to reactivity.
The oxidation states are a formal characteristic, however, appearing as one of the most important characteristics of an element.
The oxidation potential for simple substances is a measurement and indication of the potential of a substance to act in aqueous solutions, as well as the level of manifestation of redox properties.

Periodicity of elements of internal and secondary type

The periodic law gives an understanding of another important component of nature - internal and secondary periodicity. The aforementioned fields of study of atomic properties are, in fact, much more complex than one might think. This is due to the fact that the elements s, p, d of the table change their qualitative characteristics depending on their position in the period (internal periodicity) and group (secondary periodicity). For example, the internal process of the transition of the element s from the first group to the eighth to the p-element is accompanied by minimum and maximum points on the energy curve of the ionized atom. This phenomenon shows the internal inconstancy of the periodicity of changes in the properties of an atom according to its position in the period.

Results

Now the reader has a clear understanding and definition of what Mendeleev's periodic law is, realizes its significance for man and the development of various sciences, and has an idea of its current provisions and the history of discovery.

As a result of the successful development of the material in this chapter, the student should:

know

modern formulation of the periodic law;
connection between the structure of the periodic system and the energy sequence of sublevels in multielectron atoms;
definitions of the concepts "period", "group", "5-elements", "p-elements", "d- elements”, “/-elements”, “ionization energy”, “electron affinity”, “electronegativity”, “van der Waals radius”, “clarke”;
basic law of geochemistry;

be able to

Describe the structure of the periodic system in accordance with the rules of Klechkovsky;

own

Ideas about the periodic nature of the change in the properties of atoms and the chemical properties of elements, about the features of the long-period version of the periodic system; about the relationship of the abundance of chemical elements with their position in the periodic system, about macro- and microelements in the lithosphere and living matter.

Modern formulation of the periodic law

Periodic law - the most general law of chemistry - was discovered by Dmitry Ivanovich Mendeleev in 1869. At that time, the structure of the atom was not yet known. D. I. Mendeleev made his discovery based on the regular change in the properties of elements with an increase in atomic masses.

After the discovery of the structure of atoms, it became clear that their properties are determined by the structure of electron shells, which depends on total number electrons in an atom. The number of electrons in an atom is equal to the charge of its nucleus. Therefore, the modern formulation of the periodic law is as follows.

The properties of chemical elements and the simple and complex substances they form are in a periodic dependence on the charge of the nucleus of their atoms.

The significance of the periodic law lies in the fact that it is the main tool for systematizing and classifying chemical information, a very important means of interpreting chemical information, a powerful tool for predicting the properties of chemical compounds, and a means of directed search for compounds with predetermined properties.

The periodic law does not have a mathematical expression in the form of equations, it is reflected in a table called periodic system of chemical elements. There are many variants of the tables of the periodic table. The most widely used are the long-period and short-period versions, placed on the first and second color inserts of the book. The main structural unit of the periodic system is the period.

Period with number p called a sequence of chemical elements arranged in ascending order of the charge of the nucleus of an atom, which begins with ^-elements and ends with ^-elements.

In this definition P - period number equal to the main quantum number for the upper energy level in the atoms of all elements of this period. in atoms s-elements 5-sublevels are completed, in atoms p-elements - respectively p-sublevels. The exception to the above definition is the first period, in which there are no p-elements, since at the first energy level (n = 1) there is only 15-level. The periodic table also contains d-elements, whose ^-sublevels are completed, and /-elements, whose /-sublevels are completed.

Data on the structure of the nucleus and on the distribution of electrons in atoms make it possible to consider the periodic law and the periodic system of elements from fundamental physical positions. Based on modern ideas, the periodic law is formulated as follows:

The properties of simple substances, as well as the forms and properties of compounds of elements, are in a periodic dependence on the charge of the atomic nucleus (serial number).

Periodic table of D.I. Mendeleev

Currently, more than 500 variants of the representation of the periodic system are known: these are various forms of the transmission of the periodic law.

The first version of the system of elements, proposed by D.I. Mendeleev on March 1, 1869, was the so-called long form version. In this variant, the periods were arranged in one line.

In the periodic system, there are 7 horizontal periods, of which the first three are called small, and the rest are large. In the first period there are 2 elements, in the second and third - 8 each, in the fourth and fifth - 18 each, in the sixth - 32, in the seventh (incomplete) - 21 elements. Each period, with the exception of the first, begins with an alkali metal and ends with a noble gas (the 7th period is unfinished).

All elements of the periodic system are numbered in the order in which they follow each other. The element numbers are called ordinal or atomic numbers.

The system has 10 rows. Each small period consists of one row, each large period consists of two rows: even (upper) and odd (lower). In even rows of large periods (fourth, sixth, eighth and tenth) there are only metals, and the properties of the elements in the row from left to right change slightly. In odd rows of large periods (fifth, seventh and ninth), the properties of the elements in the row from left to right change, as in typical elements.

The main feature by which the elements of large periods are divided into two rows is their oxidation state. Their identical values are repeated twice in a period with an increase in the atomic masses of the elements. For example, in the fourth period, the oxidation states of elements from K to Mn change from +1 to +7, followed by the triad Fe, Co, Ni (these are elements of an even series), after which the same increase in the oxidation states of elements from Cu to Br is observed ( are elements of an odd row). We see the same in the other large periods, except for the seventh, which consists of one (even) series. The forms of combinations of elements are also repeated twice in large periods.

In the sixth period, after lanthanum, there are 14 elements with serial numbers 58-71, called lanthanides (the word "lanthanides" means similar to lanthanum, and "actinides" - "like actinium"). Sometimes they are called lanthanides and actinides, which means following lanthanide, following actinium). The lanthanides are placed separately at the bottom of the table, and in the cell an asterisk indicates the sequence of their location in the system: La-Lu. The chemical properties of the lanthanides are very similar. For example, they are all reactive metals, react with water to form Hydroxide and Hydrogen From this it follows that the lanthanides have a strong horizontal analogy.

In the seventh period, 14 elements with serial numbers 90-103 make up the actinide family. They are also placed separately - under the lanthanides, and in the corresponding cell two asterisks indicate the sequence of their location in the system: Ac-Lr. However, in contrast to the lanthanides, the horizontal analogy for actinides is weakly expressed. They exhibit more different oxidation states in their compounds. For example, the oxidation state of actinium is +3, and uranium is +3, +4, +5 and +6. The study of the chemical properties of actinides is extremely difficult due to the instability of their nuclei.

In the periodic table, eight groups are arranged vertically (indicated by Roman numerals). The group number is related to the degree of oxidation of the elements that they exhibit in compounds. As a rule, the highest positive oxidation state of elements is equal to the group number. The exceptions are fluorine - its oxidation state is -1; copper, silver, gold show oxidation states +1, +2 and +3; of the elements of group VIII, the oxidation state +8 is known only for osmium, ruthenium and xenon.

Group VIII contains the noble gases. Previously, it was believed that they are not able to form chemical compounds.

Each group is divided into two subgroups - main and secondary, which in the periodic system is emphasized by the shift of some to the right and others to the left. The main subgroup consists of typical elements (elements of the second and third periods) and elements of large periods similar to them in chemical properties. A secondary subgroup consists only of metals - elements of large periods. Group VIII is different from the others. In addition to the main helium subgroup, it contains three side subgroups: an iron subgroup, a cobalt subgroup and a nickel subgroup.

The chemical properties of the elements of the main and secondary subgroups differ significantly. For example, in VII group the main subgroup is made up of non-metals F, CI, Br, I, At, the secondary - metals Mn, Tc, Re. Thus, subgroups unite the most similar elements to each other.

All elements except helium, neon and argon form oxygen compounds; There are only 8 forms of oxygen compounds. In the periodic system, they are often represented by general formulas located under each group in ascending order of the oxidation state of the elements: R 2 O, RO, R 2 O 3, RO 2, R 2 O 5, RO 3, R 2 O 7, RO 4, where R is an element of this group. Formulas of higher oxides apply to all elements of the group (main and secondary), except for those cases when the elements do not show an oxidation state equal to the group number.

Elements of the main subgroups, starting from group IV, form gaseous hydrogen compounds, there are 4 forms of such compounds. They are also represented by general formulas in the sequence RN 4, RN 3, RN 2, RN. The formulas of hydrogen compounds are located under the elements of the main subgroups and only apply to them.

The properties of elements in subgroups change naturally: from top to bottom, metallic properties increase and non-metallic ones weaken. Obviously, the metallic properties are most pronounced in francium, then in cesium; non-metallic - in fluorine, then - in oxygen.

It is also possible to visually trace the periodicity of the properties of elements based on the consideration of the electronic configurations of atoms.

The number of electrons located at the outer level in the atoms of elements, arranged in order of increasing serial number, is periodically repeated. The periodic change in the properties of elements with an increase in the serial number is explained by the periodic change in the structure of their atoms, namely the number of electrons in their external energy levels. According to the number of energy levels in the electron shell of the atom, the elements are divided into seven periods. The first period consists of atoms in which the electron shell consists of one energy level, in the second period - of two, in the third - of three, in the fourth - of four, etc. Each new period begins when a new energy level begins to fill level.

In the periodic system, each period begins with elements whose atoms have one electron at the outer level - alkali metal atoms - and ends with elements whose atoms at the outer level have 2 (in the first period) or 8 electrons (in all subsequent ones) - noble gas atoms .

Further, we see that the outer electron shells are similar for the atoms of the elements (Li, Na, K, Rb, Cs); (Be, Mg, Ca, Sr); (F, Cl, Br, I); (He, Ne, Ag, Kr, Xe), etc. That is why each of the above groups of elements is in a certain main subgroup of the periodic table: Li, Na, K, Rb, Cs in group I, F, Cl, Br, I - in VII, etc.

It is precisely because of the similarity of the structure of the electron shells of atoms that their physical and chemical properties are similar.

Number main subgroups is determined by the maximum number of elements at the energy level and is equal to 8. The number of transition elements (elements side subgroups) is determined by the maximum number of electrons in the d-sublevel and is equal to 10 in each of the large periods.

Since in the periodic system of chemical elements D.I. Mendeleev, one of the side subgroups contains at once three transition elements that are close in chemical properties (the so-called Fe-Co-Ni, Ru-Rh-Pd, Os-Ir-Pt triads), then the number of side subgroups, as well as the main ones, is eight.

By analogy with the transition elements, the number of lanthanides and actinides placed at the bottom of the periodic system in the form of independent rows is equal to the maximum number of electrons at the f-sublevel, i.e. 14.

The period begins with an element in the atom of which there is one s-electron at the outer level: in the first period it is hydrogen, in the rest - alkali metals. The period ends with a noble gas: the first - with helium (1s 2), the remaining periods - with elements whose atoms at the outer level have an electronic configuration ns 2 np 6 .

The first period contains two elements: hydrogen (Z = 1) and helium (Z = 2). The second period begins with the element lithium (Z= 3) and ends with neon (Z= 10). There are eight elements in the second period. The third period begins with sodium (Z = 11), the electronic configuration of which is 1s 2 2s 2 2p 6 3s 1. The filling of the third energy level began from it. It ends at the inert gas argon (Z= 18), whose 3s and 3p sublevels are completely filled. Electronic formula of argon: 1s 2 2s 2 2p 6 Zs 2 3p 6. Sodium is an analogue of lithium, argon is an analogue of neon. In the third period, as in the second, there are eight elements.

The fourth period begins with potassium (Z = 19), the electronic structure of which is expressed by the formula 1s 2 2s 2 2p 6 3s 2 3p64s 1. Its 19th electron occupied the 4s sublevel, the energy of which is lower than the energy of the 3d sublevel. The outer 4s electron gives the element properties similar to those of sodium. In calcium (Z = 20), the 4s sublevel is filled with two electrons: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2. From the scandium element (Z = 21), the filling of the 3d sublevel begins, since it is energetically more favorable than 4p -sublevel. Five orbitals of the 3d sublevel can be occupied by ten electrons, which occurs in atoms from scandium to zinc (Z = 30). Therefore, the electronic structure of Sc corresponds to the formula 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 4s 2, and zinc - 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2. In the atoms of subsequent elements up to the inert gas krypton (Z = 36) the 4p sublevel is being filled. There are 18 elements in the fourth period.

The fifth period contains elements from rubidium (Z = 37) to the inert gas xenon (Z = 54). The filling of their energy levels is the same as for the elements of the fourth period: after Rb and Sr, ten elements from yttrium (Z= 39) to cadmium (Z = 48), the 4d sublevel is filled, after which the electrons occupy the 5p sublevel. In the fifth period, as in the fourth, there are 18 elements.

In atoms of elements of the sixth period of cesium (Z= 55) and barium (Z = 56), the 6s sublevel is filled. In lanthanum (Z = 57), one electron enters the 5d sublevel, after which the filling of this sublevel stops, and the 4f sublevel begins to fill, seven orbitals of which can be occupied by 14 electrons. This occurs for atoms of the lanthanide elements with Z = 58 - 71. Since these elements fill the deep 4f sublevel of the third level from the outside, they have very similar chemical properties. With hafnium (Z = 72), the filling of the d-sublevel resumes and ends with mercury (Z = 80), after which the electrons fill the 6p-sublevel. The filling of the level is completed at the noble gas radon (Z = 86). There are 32 elements in the sixth period.

The seventh period is incomplete. The filling of electronic levels with electrons is similar to the sixth period. After filling the 7s sublevel in France (Z = 87) and radium (Z = 88), an actinium electron enters the 6d sublevel, after which the 5f sublevel begins to be filled with 14 electrons. This occurs for atoms of actinide elements with Z = 90 - 103. After the 103rd element, the b d-sublevel is filled: in kurchatovium (Z = 104), = 105), elements Z = 106 and Z = 107. Actinides, like lanthanides, have many similar chemical properties.

Although the 3d sublevel is filled after the 4s sublevel, it is placed earlier in the formula, since all sublevels of this level are written sequentially.

Depending on which sublevel is last filled with electrons, all elements are divided into four types (families).

1. s - Elements: the s-sublevel of the outer level is filled with electrons. These include the first two elements of each period.

2. p - Elements: the p-sublevel of the outer level is filled with electrons. These are the last 6 elements of each period (except the first and seventh).

3. d - Elements: the d-sublevel of the second level from the outside is filled with electrons, and one or two electrons remain at the outer level (for Pd - zero). These include elements of intercalary decades of large periods located between s- and p-elements (they are also called transitional elements).

4. f - Elements: the f-sublevel of the third level from the outside is filled with electrons, and two electrons remain at the outer level. These are the lanthanides and actinides.

There are 14 s-elements, 30 p-elements, 35 d-elements, 28 f-elements in the periodic system. Elements of the same type have a number of common chemical properties.

The periodic system of D. I. Mendeleev is a natural classification of chemical elements according to the electron structure of their atoms. The electronic structure of an atom, and hence the properties of an element, is judged by the position of the element in the corresponding period and subgroup of the periodic system. The patterns of filling of electronic levels explain the different number of elements in periods.

Thus, the strict periodicity of the arrangement of elements in the periodic system of chemical elements of D. I. Mendeleev is fully explained by the consistent nature of the filling of energy levels.

Findings:

The theory of the structure of atoms explains the periodic change in the properties of elements. An increase in the positive charges of atomic nuclei from 1 to 107 causes a periodic repetition of the structure of the external energy level. And since the properties of the elements mainly depend on the number of electrons in the outer level, they also repeat periodically. This is the physical meaning of the periodic law.

In short periods, with an increase in the positive charge of the nuclei of atoms, the number of electrons at the external level increases (from 1 to 2 - in the first period, and from 1 to 8 - in the second and third periods), which explains the change in the properties of the elements: at the beginning of the period (except for the first period) there is an alkali metal, then the metallic properties gradually weaken and the non-metallic properties increase.

In large periods, as the nuclear charge increases, filling the levels with electrons is more difficult, which also explains the more complex change in the properties of elements compared to elements of small periods. So, in even rows of long periods, with increasing charge, the number of electrons in the outer level remains constant and is equal to 2 or 1. Therefore, while the electrons are filling the level following the outer (second from the outside), the properties of the elements in these rows change extremely slowly. Only in odd rows, when the number of electrons in the outer level increases with the growth of the nuclear charge (from 1 to 8), do the properties of the elements begin to change in the same way as for typical ones.

In the light of the doctrine of the structure of atoms, the division of D.I. Mendeleev of all elements for seven periods. The period number corresponds to the number of energy levels of atoms filled with electrons. Therefore, s-elements are present in all periods, p-elements in the second and subsequent, d-elements in the fourth and subsequent, and f-elements in the sixth and seventh periods.

The division of groups into subgroups, based on the difference in the filling of energy levels with electrons, is also easily explained. For elements of the main subgroups, either s-sublevels (these are s-elements) or p-sublevels (these are p-elements) of the outer levels are filled. For elements of side subgroups, the (d-sublevel of the second outside level (these are d-elements) is filled. For lanthanides and actinides, the 4f- and 5f-sublevels are filled, respectively (these are f-elements). Thus, in each subgroup, elements are combined whose atoms have similar structure of the outer electronic level.At the same time, the atoms of the elements of the main subgroups contain on the outer levels the number of electrons equal to the number of the group.The secondary subgroups include elements whose atoms have on the outer level two or one electron.

Differences in structure also cause differences in the properties of elements of different subgroups of the same group. So, at the outer level of the atoms of the elements of the halogen subgroup, there are seven electrons of the manganese subgroup - two electrons each. The former are typical metals and the latter are metals.

But the elements of these subgroups also have common properties: entering into chemical reactions, all of them (with the exception of fluorine F) can donate 7 electrons to form chemical bonds. In this case, the atoms of the manganese subgroup donate 2 electrons from the outer and 5 electrons from the next level. Thus, in the elements of the secondary subgroups, the valence electrons are not only the outer, but also the penultimate (second from the outside) levels, which is the main difference in the properties of the elements of the main and secondary subgroups.

It also follows that the group number, as a rule, indicates the number of electrons that can participate in the formation of chemical bonds. This is the physical meaning of the group number.

So, the structure of atoms determines two patterns:

1) change in the properties of elements horizontally - in the period from left to right, metallic properties are weakened and non-metallic properties are enhanced;

2) a change in the properties of elements along the vertical - in a subgroup with an increase in the serial number, metallic properties increase and non-metallic ones weaken.

In this case, the element (and the cell of the system) is located at the intersection of the horizontal and vertical, which determines its properties. This helps to find and describe the properties of elements whose isotopes are obtained artificially.

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 courses in chemistry, physics and biology.

Mendeleev's periodic law is one of the fundamental laws of chemistry. It can be argued that all modern chemistry is built on it. He explains the dependence of the properties of atoms on their structure, generalizes this dependence for all elements, dividing them into different groups, and also predicts their properties depending on the structure and structure depending on the properties.

There are other laws that have explanatory, generalizing and predictive functions. For example, the law of conservation of energy, the law of refraction of light, Mendel's genetic law.

2. Name the chemical element in whose atom the electrons are arranged in levels according to a series of numbers: 2, 5. What simple substance forms this element? What is the formula of its hydrogen compound and what is its name? What formula does the highest oxide of this element have, what is its character? Write down the reaction equations characterizing the properties of this oxide.

3. Beryllium used to be classified as a group III element, and its relative atomic mass was considered to be 13.5. Why did D. I. Mendeleev transfer it to group II and correct the atomic mass of beryllium from 13.5 to 9?

Previously, the element beryllium was mistakenly assigned to group III. The reason for this was the incorrect determination of the atomic mass of beryllium (instead of 9, it was considered equal to 13.5). D. I. Mendeleev suggested that beryllium is in group II, based on chemical properties element. The properties of beryllium were very similar to those of Mg and Ca, and completely different from those of Al. Knowing that the atomic masses of Li and B, neighboring elements to Be, are 7 and 11, respectively, D. I. Mendeleev suggested that the atomic mass of beryllium is 9.

4. Write the equations of reactions between a simple substance formed by a chemical element in the atom of which electrons are distributed over 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 system. What is the type of chemical bond in the reaction products? What is the crystalline structure of the initial simple substances and the products of their interaction?

5. Arrange the following elements in order of strengthening the metallic properties: As, Sb, N, P, Bi. Justify the resulting series based on the structure of the atoms of these elements.

N, P, As, Sb, Bi - strengthening of metallic properties. The metallic properties in the groups are enhanced.

6. Arrange the following elements in order of strengthening non-metallic properties: Si, Al, P, S, Cl, Mg, Na. Justify the resulting series based on the structure of the atoms of these elements.

Na, Mg, Al, Si, P, S, Cl - strengthening of non-metallic properties. Non-metallic properties in periods are enhanced.

7. Arrange in the order of weakening the acid properties of the oxides, the formulas of which are: SiO2, P2O5, Al2O3, Na2O, MgO, Cl2O7. Justify the resulting series. Write down the formulas of the hydroxides corresponding to these oxides. How does their acid character change in the series you proposed?

8. Write the formulas for the oxides of boron, beryllium and lithium and arrange them in ascending order of the 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 formation of the Periodic Law?

The periodic system of elements reflects the relationship of chemical elements. The atomic number of an element is equal to the charge of the nucleus, numerically it is equal to the number protons. The number of neutrons contained in the nuclei of one element, in contrast to the number of protons, can be different. Atoms of the same element whose nuclei contain different number neutrons are called isotopes.

Each chemical element has several isotopes (natural or artificial). The atomic mass of a chemical element is equal to the average value of the masses of all its natural isotopes, taking into account their abundance.

With the discovery of isotopes, the charges of nuclei, rather than their atomic masses, began to be used to distribute elements in the periodic system.

10. Why do the charges of the atomic nuclei of elements in the Periodic system of D. I. Mendeleev change monotonically, i.e., 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?

This is due to the fact that the properties of elements and their compounds do not depend on the total number of electrons, but only on the valence electrons that are on the last layer. The number of valence electrons changes periodically, therefore, the properties of elements also change periodically.

11. Give three formulations of the Periodic Law, in which the relative atomic mass, the 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.

1. The properties of chemical elements and the substances formed by them are in a periodic dependence on the relative atomic masses of the elements.
2. The properties of chemical elements and the substances formed by them are in a periodic dependence on the charge of the atomic nuclei of the elements.
3. The properties of chemical elements and the substances formed by them are in a periodic dependence on the structure of external energy levels in the electron shell of an atom.