"The properties of the elements, and therefore the simple and complex bodies (substances) formed by them, stand in a periodic dependence on their atomic weight."

Modern wording:

"the properties of chemical elements (i.e., the properties and form of the compounds they form) are in a periodic dependence on the charge of the nucleus of atoms of chemical elements."

The physical meaning of chemical periodicity

Periodic changes in the properties of chemical elements are due to the correct repetition of the electronic configuration of the external energy level (valence electrons) of their atoms with an increase in the nuclear charge.

The graphic representation of the periodic law is the periodic table. It contains 7 periods and 8 groups.

Period - horizontal rows of elements with the same maximum value of the main quantum number of valence electrons.

The period number denotes the number of energy levels in an element's atom.

Periods can consist of 2 (first), 8 (second and third), 18 (fourth and fifth), or 32 (sixth) elements, depending on the number of electrons in the outer energy level. The last, seventh period is incomplete.

All periods (except the first) begin with an alkali metal ( s- element) and end with a noble gas ( ns 2 np 6 ).

Metallic properties are considered as the ability of atoms of elements to easily donate electrons, and non-metallic properties to accept electrons due to the tendency of atoms to acquire a stable configuration with filled sublevels. Filling the outer s- sublevel indicates the metallic properties of the atom, and the formation of the outer p- sublevel - on non-metallic properties. An increase in the number of electrons by p- sublevel (from 1 to 5) enhances the non-metallic properties of the atom. Atoms with a fully formed, energetically stable configuration of the outer electron layer ( ns 2 np 6 ) chemically inert.

In long periods, the transition of properties from the active metal to the noble gas occurs more smoothly than in short periods, because the formation of an internal n - 1) d - sublevel while maintaining the external ns 2 - layer. Large periods consist of even and odd rows.

For elements of even rows on the outer layer ns 2 - electrons, therefore, metallic properties predominate and their weakening with increasing nuclear charge is small; in odd rows is formed np- sublevel, which explains the significant weakening of the metallic properties.

Groups - vertical columns of elements with the same number of valence electrons, equal to the group number. There are main and secondary subgroups.

The main subgroups consist of elements of small and large periods, the valence electrons of which are located on the outer ns - and np - sublevels.

Secondary subgroups consist of elements of only large periods. Their valence electrons are on the outer ns- sublevel and internal ( n - 1) d - sublevel (or (n - 2) f - sublevel).

Depending on which sublevel ( s-, p-, d- or f-) filled with valence electrons, the elements of the periodic system are divided into: s- elements (elements of the main subgroup I and II groups), p - elements (elements of the main subgroups III - VII groups), d - elements (elements of secondary subgroups), f- elements (lanthanides, actinides).

In the main subgroups, from top to bottom, metallic properties are enhanced, while non-metallic properties are weakened. The elements of the main and secondary groups differ greatly in properties.

The group number indicates the highest valency of the element (except O , F , elements of the copper subgroup and the eighth group).

Common to the elements of the main and secondary subgroups are the formulas of higher oxides (and their hydrates). For higher oxides and their element hydrates I-III groups (except for boron) the basic properties predominate, with IV to VIII - acidic.

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The periodic repetition of the properties of elements with increasing atomic number becomes especially evident if the elements are arranged in the form of a table, called the periodic table or the periodic system of elements. Several forms of the periodic table have been proposed and are in use.

The periodic repetition of the properties of the elements with increasing atomic number can be clearly shown if the elements are arranged in a table called the periodic table, or the periodic system, of the elements. Many different forms of the periodic table have been proposed and are in use.

The principle of periodic repetition of the properties of elements could not allow the existence of only one, isolated element of argon; such simple substances should be few or none. However, Ramsay firmly stood on the positions of the periodic law, and this, as well as the development of laboratory technology at the end of the last century, predetermined the rapid discovery of the remaining members of the group of inert gases.

What explains the periodic repetition of the properties of elements in the periodic system.

What explains the periodic repetition of the properties of elements.

Accepting that the periodic repetition of the properties of elements is due not only to their mass (atomic weight), but also to the nature of the movement of the atoms themselves as whole particles (the speed and direction of their movement), Flavitsky builds his hypothesis on the following basis: the periodicity of elements is explained not by what is repeated type of the internal structure of atoms, but by the fact that the nature of the movement of atoms as whole particles periodically changes.

Thus, the reason for the periodic repetition of the properties of elements is the periodic repetition of the electronic configurations of their atoms.

The study of the electronic structure of atoms made it possible to prove that the reason for the periodic repetition of the properties of elements with increasing serial number is the periodic repetition of the process of building new electron shells. To the same group of the periodic system always belong those elements whose atoms in the outer shells have the same number of electrons. So, the atoms of all inert gases, except for helium, contain 8 electrons in the outer shell and are the most difficult to ionize, while the atoms of alkali metals contain one electron in the outer shell and have the lowest ionization potential. Alkali metals with only one electron in the outer shell can easily lose it, turning into a stable form of a positive ion with an electron configuration similar to the nearest inert gas with a lower atomic number. Elements such as fluorine, chlorine, etc., approaching the configuration of inert gases in terms of the number of external electrons, on the contrary, tend to acquire electrons and reproduce this electronic configuration, passing into the corresponding negative ion.

The periods following the third period of the table of D. I. Mendeleev are longer. However, the periodic repetition of the properties of the elements is preserved. It becomes more complex, due to the increasing variety of physical and chemical features of the elements as their atomic masses increase. Consideration of the structure of atoms of the first periods confirms that the limited number of places for electrons in each shell (Pauli prohibition) surrounding the nucleus is the reason for the periodic repetition of the properties of elements. This periodicity is a great law of nature, discovered by D. I. Mendeleev at the end of the last century, in our time has become one of the foundations for the development of not only chemistry, but also physics.

The values ​​of /j increase gradually as Z increases until Z reaches the noble gas value and then drops to about one-fourth the noble gas value as it moves to the next element. The periodicity of changes in another property - the density of elements in the solid state - is shown in fig. 5.13. Such a periodic repetition of the properties of elements with an increase in the serial number becomes especially evident if the elements are arranged in the form of a table called the periodic table and the periodic system of elements. Many different forms of the periodic system have been proposed and are in use.

Simultaneously with Newlands, de Chancourtois was approaching the discovery of the periodic law in France. But in contrast to the sensual musical and sound image, which served for Newlands as an analogy with the regularity of chemical elements that he partially revealed, the French naturalist used an abstract geometric image: he compared the periodic repetition of the properties of the elements, arranged according to their atomic weights, with the winding of a spiral line (vis tellurique) and the side surface of the cylinder.

The idea of ​​the magnitude of the charge of the nucleus as the defining property of the atom formed the basis of the modern formulation of the periodic law of D. I. Mendeleev: the properties of chemical elements, as well as the forms and properties of the compounds of these elements, are in a periodic dependence on the magnitude of the charge of the nuclei of their atoms. It made it possible to explain the reason for the periodic repetition of the properties of elements, which consists in the periodic repetition of the structure of the electronic configurations of atoms.

Only after the structure of the atom was clarified did the reasons for the periodic repetition of the properties of elements become clear.

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 group VII, the main subgroup is made up of non-metals F, CI, Br, I, At, while the side group is 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 the deep 4f sublevel of the third level from the outside is filled in these elements, 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: when 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.

Periodicity in changing the properties of elements. Periodic law D.I. Mendeleev

The periodic system of chemical elements was created in 1869 by our great compatriot Dmitry Ivanovich Mendeleev.

Unlike his predecessors, Mendeleev compared not only similar, but mostly dissimilar elements and their groups (for example, alkali metals and halogens), placing them on the basis of the main (known by that time) characteristic of the element - atomic weight.

The wording of the law at the time was:

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

Later, Mendeleev used the characteristic of elements that he introduced, more fundamental than atomic weight, namely their serial number, which is determined by the positive charge of the nucleus, i.e. the number of protons in the nucleus of an atom. Regularities were established for changing the properties of elements in periods and groups.

To describe and systematize chemical elements, it is necessary to know their characteristics: serial number (charge of the nucleus of its atoms) and relative atomic mass.

Of these, the charge of the nucleus of atoms is a common, unchanged during chemical reactions, the main characteristic for determining the element.

To describe the elements, in addition to the quantitative characteristics listed above, others are needed, including the qualitative characteristics of the element. These are the electronic structure and properties of its atoms.

Of particular importance are the electrons located on the outer electron layer, the valence electrons. For metal elements, they usually have 1 - 2, less often 3, for non-metals - 4 or more. For elements of large periods of side subgroups, the valence electrons are not only the outer, but also the pre-outer layer. The reactivity of atoms to form chemical bonds with other atoms, to form chemical compounds, depends on valence electrons.

A chemical compound is a chemically individual substance, consisting of chemically bonded atoms of one element in a simple or several elements in a complex substance, having a certain composition.

Simple and complex substances are forms of the real existence of elements in nature. The nature of the elements affects the properties of the substances formed by them, and vice versa, knowing the properties of substances, one can judge the nature of the element.

Dmitri Ivanovich Mendeleev attached great importance to the knowledge of the forms and properties of typical oxygen and hydrogen compounds of an element for its characterization. Under the form of compounds, he understood the similarity in the composition of their compounds typical for a group of elements, expressed by general formulas. Thus, the elements of the main subgroup of group VI of the periodic system have the following forms of oxygen and hydrogen compounds: RO3, H2R.

For example: sulfur oxide and hydrogen sulfide.

Typical metallic elements form basic oxides and hydroxides, exhibiting low valence values ​​in these forms of compounds. In non-metallic elements, higher oxygen compounds (oxides and hydroxides) are acidic. These elements form gaseous hydrogen compounds. Many elements exhibit intermediate properties.

Let us derive the patterns of changes in the properties of elements with an increase in their serial number.

1. The most important quantitative characteristics of an element - the charge of the nucleus of its atoms and the atomic mass - increase monotonically.

2. The structures of the outer electronic layer change abruptly.

3. Forms and properties of oxides and hydroxides of elements are periodically repeated.

4. Periodically, the valence of elements in oxygen increases and decreases in hydrogen.

What is the relationship between the characteristics of the element, changing monotonically and periodically?

Let us consider this relationship using the example of the charge of the nucleus of atoms and their outer electrons. To do this, we will build a graph. Note on the horizontal line the charge of the atomic nucleus, and on the vertical line - the number of electrons in the outer layer of the atoms of the elements.

The number of electrons in the outer electron layer of the atoms of elements periodically changes with a monotonous increase in the charge of the nucleus of their atoms.

The discovery of the periodic law marked the beginning of a new era in the development of chemistry - its modern stage. Prior to this, the facts accumulated in science had no internal connection.

The periodic law revealed a deep connection between the elements, allowed scientists to predict the properties of yet undiscovered elements and their compounds, and purposefully search for new ones.

Dmitri Ivanovich Mendeleev did not doubt the reliability of the open law, he firmly believed in its future, in its development. Shortly before his death, he wrote: "... the future does not threaten the periodic law with destruction, but only promises superstructures and development."

Periodic law:

Approved a deep internal connection between the elements;

Allowed scientists to assume that all atoms are built according to a common plan;

Thus, he created a prerequisite for the transition to a new stage in the development of science, to the knowledge of the internal structure of atoms - the discovery of the electron, radioactivity, the development of a theory of the structure of the atom, etc.

The next step was the disclosure of the physical essence of the law based on the theory of the structure of the atom.

You are already familiar with the structure of atoms and you know that the charge of the nucleus of an atom is its main characteristic. The charge of the nucleus coincides with the ordinal number of the element in the periodic system of Mendeleev.

Rutherford's student, the English physicist Henry Moseley, established in 1913 that each element has its own wavelength of X-ray radiation. It increases with increasing atomic mass. Moseley related the frequency of this radiation to the ordinal number of the element. Moseley's law confirmed that Mendeleev's change in the serial numbers of elements in the periodic system corresponded to a consistent increase in the charges of the nuclei of their atoms. We have already discussed this question in the study of isotopes.

In connection with new discoveries in the field of atomic structure, the periodic law adopted the following modern formulation:

The properties of the elements, as well as the forms and properties of their compounds, are in a periodic dependence on the charge of the atomic nucleus.

Why do the properties of elements and their compounds change periodically?

What is the reason for the periodicity?

The answer to this question can also be given by the theory of the structure of the atom:

The value of the charge of the nucleus is the main characteristic of the element, a measure of its individuality. All other properties of the element depend on this characteristic of the element; it determines the number of electrons and their state in the atom.

The increase in the charges of atomic nuclei from the first to the last element leads to a periodic repetition of the electronic structures of atoms and the number of electrons at the external energy level. This is the physical meaning of the periodic law and the reason for the periodicity of changes in the properties of elements.

The periodic change in the properties of elements is explained by the periodic repetition of the number of electrons at the external energy level and the electronic structures of atoms.

The theory of the structure of the atom contributed to the development of the periodic law and the periodic system of chemical elements, the determination of their modern content. It gave impetus to the study of the internal structure of substances, to the discovery and production of new elements.

The charges of the nuclei of elements in the periodic system are continuously increasing, and the properties of simple substances are repeated periodically. How to explain it?

D. I. Mendeleev noticed that the properties of elements are periodically repeated with increasing values ​​of their mass numbers. He arranged the 63 elements discovered by that time in order of increasing their atomic masses, taking into account chemical and physical properties. Mendeleev believed that the periodic law he discovered was a reflection of deep patterns in the internal structure of matter, he stated the fact of periodic changes in the properties of elements, but did not know the reason for the periodicity.

Further study of the structure of the atom showed that the properties of substances depend on the charge of the nucleus of atoms, and elements can be systematized based on their electronic structure. The properties of simple substances and their compounds depend on the periodically repeating electronic configuration of the valence sublevel of the element's atoms. Therefore, "electronic analogues" are also "chemical analogues".

Let us write down the electronic formulas of the atoms of the elements of the main subgroups of the second and seventh groups.

Elements of the second group have the general electronic formula of valence electrons ns 2 . Let's write down their electronic formulas:

Be 1s 2 2s 2,

Mg 1s 2 2s 2 2p 6 3s 2,

Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2,

Sr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2.

The elements of the seventh group have a common electronic formula of valence electrons ns 2 np 5, and the full electronic formulas look like:

F 1s 2 2s 2 2p 5 ,

Cl 1s 2 2s 2 2p 6 3s 2 3p 5 ,

Br 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p5 ,

I 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p5 .

So, the electronic structures of atoms are periodically repeated for elements of the same group, therefore their properties are periodically repeated, since they depend mainly on the electronic configuration of valence electrons. Elements of the same group have common properties, but there are also differences. This can be explained by the fact that although atoms have the same electronic structure of valence electrons, these electrons are located at different distances from the nucleus, the force of their attraction to the nucleus weakens during the transition from period to period, the atomic radius increases, valence electrons become more mobile, which affects the properties of substances.

41. Based on the position of germanium, cesium and technetium in the periodic system, make formulas for the following compounds: meta and orthogermanic acids, cesium dihydrogen phosphate and technetium oxide, corresponding to its highest oxidation state. Draw the structural formulas of these compounds.

42. What is ionization energy? In what units is it expressed? How does the reducing activity of s- and p-elements in the groups of the periodic system change with increasing serial number? Why?

43. What is electronegativity? How does the electronegativity of elements in the second and third periods, in the group of the periodic system, change with increasing serial number?

44. Based on the position of germanium, molybdenum and rhenium in the periodic system, make up the gross formulas of the following compounds: the hydrogen compound of germanium, rhenium acid and molybdenum oxide, corresponding to its highest oxidation state. Draw the structural formulas of these compounds.

45. What is electron affinity? In what units is it expressed? How does the oxidative activity of non-metals change in a period and in a group of the periodic system with an increase in the serial number? Justify your answer by the structure of the atom of the corresponding element.

46. ​​Make formulas for oxides and hydroxides of elements of the third period of the periodic system, corresponding to their highest oxidation state. How does the chemical nature of these compounds change when going from sodium to chlorine?

47. Which of the elements of the fourth period - vanadium or arsenic - has more pronounced metallic properties? Which element forms a gaseous compound with hydrogen? Justify your answer based on the structure of the atoms of these elements.

48. What elements form gaseous compounds with hydrogen? What groups of the periodic table are these elements in? Write formulas for hydrogen and oxygen compounds of chlorine, tellurium and antimony corresponding to their lowest and highest oxidation states.

49. Which element of the fourth period - chromium or selenium - has more pronounced metallic properties? Which of these elements forms a gaseous compound with hydrogen? Motivate your answer by the structure of the atoms of chromium and selenium.

50. What is the lowest oxidation state of chlorine, sulfur, nitrogen and carbon? Why? Write formulas for aluminum compounds with these elements in their oxidation state. What are the names of the corresponding compounds?

51. Which of the p-elements of the fifth group of the periodic system - phosphorus or antimony - has non-metallic properties more pronounced? Which of the hydrogen compounds of these elements is the stronger reducing agent? Justify your answer by the structure of the atom of these elements.

52. Based on the position of the metal in the periodic system, give a reasoned answer to the question; which of the two hydroxides is the stronger base: Ba(OH) 2 or Mg(OH) 2; Ca(OH) 2 or Fe(OH) 2; Cd (OH) 2 or Sr (OH) 2?

53. Why does manganese exhibit metallic properties, and chlorine non-metallic? Motivate your answer by the electronic structure of the atoms of these elements. Write the formulas for oxides and hydroxides of chlorine and manganese.

54. What is the lowest oxidation state of hydrogen, fluorine, sulfur and nitrogen? Why? Write formulas for calcium compounds with these elements in their oxidation state. What are the names of the corresponding compounds?

55. What are the lowest and highest oxidation states of silicon, arsenic, selenium and chlorine? Why? Write formulas for compounds of these elements corresponding to these oxidation states.

56. To which family do the elements belong, in the atoms of which the last electron enters the 4f- and 5f-orbitals? How many elements does each of these families include?

57. The atomic masses of the elements in the periodic system are continuously increasing, while the properties of simple bodies change periodically. How can this be explained?

58. What is the modern formulation of the periodic law? Explain why in the periodic table of elements argon, cobalt, tellurium and thorium are placed before potassium, nickel, iodine and protactinium, respectively, although they have a large atomic mass?

59. What are the lowest and highest oxidation states of carbon, phosphorus, sulfur and iodine? Why? Write formulas for compounds of these elements corresponding to these oxidation states.