Regularities of the periodic table of chemical elements. The properties of elements are periodically dependent on the charge of their atomic nuclei

The periodic law and the periodic system of chemical elements of D. I. Mendeleev based on ideas about the structure of atoms. The importance of the periodic law for the development of science.

In 1869, D.I. Mendeleev, based on an analysis of the properties of simple substances and compounds, formulated the Periodic Law:

The properties of simple bodies... and compounds of elements are periodically dependent on the magnitude of the atomic masses of the elements.

Based on the periodic law, the periodic system of elements was compiled. In it, elements with similar properties were combined into vertical columns - groups. In some cases, when placing elements in the Periodic Table, it was necessary to disrupt the sequence of increasing atomic masses in order to maintain the periodicity of the repetition of properties. For example, we had to “swap” tellurium and iodine, as well as argon and potassium.

The reason is that Mendeleev proposed the periodic law at a time when nothing was known about the structure of the atom.

After the planetary model of the atom was proposed in the 20th century, the periodic law was formulated as follows:

The properties of chemical elements and compounds periodically depend on the charges of atomic nuclei.

The charge of the nucleus is equal to the number of the element in the periodic table and the number of electrons in the electron shell of the atom.

This formulation explained the "violations" of the Periodic Law.

In the Periodic Table, the period number is equal to the number of electronic levels in the atom, the group number for elements of the main subgroups is equal to the number of electrons in the outer level.

The reason for the periodic change in the properties of chemical elements is the periodic filling of electron shells. After filling the next shell, a new period begins. The periodic change of elements is clearly visible in the changes in the composition and properties of the oxides.

Scientific significance of the periodic law. The periodic law made it possible to systematize the properties of chemical elements and their compounds. When compiling the periodic table, Mendeleev predicted the existence of many undiscovered elements, leaving empty cells for them, and predicted many properties of undiscovered elements, which facilitated their discovery

Ticket No. 2

The structure of atoms of chemical elements using the example of elements of the second period and IV-A group of the periodic system of chemical elements by D. I. Mendeleev. Regularities in the changes in the properties of these chemical elements and the simple and complex substances formed by them (oxides, hydroxides) depending on the structure of their atoms.

As you move from left to right along a period, the metallic properties of the elements become less and less pronounced. When moving from top to bottom within one group, elements, on the contrary, display increasingly pronounced metallic properties. Elements located in the middle part of the short periods (2nd and 3rd periods) usually have a skeleton covalent structure, and elements from the right part of these periods exist in the form of simple covalent molecules.

Atomic radii change as follows: decrease when moving from left to right along a period; increase as you move from top to bottom along the group. As you move from left to right across a period, electronegativity, ionization energy, and electron affinity increase, reaching a maximum for the halogens. For noble gases, electronegativity is 0. Changes in the electron affinities of elements when moving from top to bottom along the group are not so characteristic, but at the same time the electronegativity of the elements decreases.

In elements of the second period, 2s and then 2p orbitals are filled.

The main subgroup of group IV of the periodic system of chemical elements by D. M. Mendeleev contains carbon C, silicon Si, germanium Ge, tin Sn and lead Pb. The outer electron layer of these elements contains 4 electrons (s 2 p 2 configuration). Therefore, the elements of the carbon subgroup must have some similarities. In particular, their highest oxidation state is the same and is +4.

What causes the difference in the properties of the elements of the subgroup? The difference between the ionization energy and the radius of their atoms. As the atomic number increases, the properties of elements naturally change. Thus, carbon and silicon are typical non-metals, tin and lead are metals. This is manifested primarily in the fact that carbon forms a simple non-metal substance (diamond), and lead is a typical metal.

Germanium occupies an intermediate position. According to the structure of the electronic shell of the atom, p-elements of group IV have even oxidation states: +4, +2, – 4. The formula of the simplest hydrogen compounds is EN 4, and the E-H bonds are covalent and equivalent due to the hybridization of s- and p-orbitals with the formation sp 3 orbitals directed at tetrahedral angles.

The weakening of the characteristics of a non-metallic element means that in the subgroup (C-Si-Ge-Sn-Pb) the highest positive oxidation state +4 becomes less and less characteristic, and the oxidation state +2 becomes more typical. So, if for carbon the most stable compounds are those in which it has an oxidation state of +4, then for lead the compounds in which it exhibits an oxidation state of +2 are most stable.

What can be said about the stability of compounds of elements in the negative oxidation state -4? Compared to non-metallic elements of groups VII-V, p-elements of group IV exhibit signs of a non-metallic element to a lesser extent. Therefore, for elements of the carbon subgroup, a negative oxidation state is atypical.

DI. Mendeleev formulated the Periodic Law in 1869, which was based on one of the most important characteristics of an 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 now, there are 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 structures and which are 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 the distribution of electrons over the outer atomic orbital plays a major role. 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 population). In short periods with increasing nuclear charge population 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

2. does not exceed 2, and then up to 8. 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. By period, the atomic radius decreases (gradually adding electrons are described by orbitals with almost equal characteristics; by group, the atomic radius increases as the number of electron layers increases (Fig. 1).

Rice. 1. Periodic change in atomic radius

3. 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). 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< Е и 2 < Е и 3 <….Энергии ионизации отражают дискретность структуры электронных слоев и оболочек атомов химических элементов.

4. 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 by reporting discrete values ​​of E and. Moreover, E and 1

Electron affinity

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

Patterns of changes in the chemical properties of elements and their compounds by periods and groups

Let us list the patterns of changes in properties that appear within periods:

— metallic properties decrease;

— non-metallic properties are enhanced;

— the degree of oxidation of elements in higher oxides increases from $+1$ to $+7$ ($+8$ for $Os$ and $Ru$);

— the degree of oxidation of elements in volatile hydrogen compounds increases from $-4$ to $-1$;

- oxides from basic through amphoteric are replaced by acidic oxides;

- hydroxides from alkalis through amphoteric ones are replaced by acids.

D.I. Mendeleev made a conclusion in $1869 - he formulated the Periodic Law, which sounds like this:

The properties of chemical elements and the substances formed by them are periodically dependent on the relative atomic masses of the elements.

Systematizing chemical elements based on their relative atomic masses, Mendeleev also paid great attention to the properties of the elements and the substances they form, distributing elements with similar properties into vertical columns - groups.

Sometimes, in violation of the pattern he discovered, Mendeleev placed heavier 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. Mendeleev considered this order of arrangement necessary because otherwise these elements would fall into groups of elements dissimilar to them in properties, 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, nor could he explain the reason for the periodicity of the properties of elements and the substances formed by them. However, he foresaw that this reason lay in the complex structure of the atom, the internal structure of which was not studied at that time.

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 properties of chemical elements and the substances formed by them are periodically dependent on the charges of their atomic nuclei.

The periodicity in changes in the properties of elements 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 Table, i.e. reveal the physical meaning of the period number, group number and ordinal number of the element.

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.

The periodic law and the Periodic System of Chemical Elements by 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 are complemented by one more, which is the ability to predict, i.e. predict, describe properties and indicate ways to discover new chemical elements.

General characteristics of metals of the main subgroups of groups I±III in connection with their position in the Periodic Table of Chemical Elements of D. I. Mendeleev and the structural features of their atoms

Chemical elements - metals

Most chemical elements are classified as metals—$92 of the $114 known elements.

All metals, except mercury, in their normal state are solids and have a number of common properties.

Metals- These are malleable, plastic, viscous substances that have a metallic luster and are capable of conducting heat and electric current.

Atoms of metal elements give up electrons from the outer (and some from the outer) electron layer, turning into positive ions.

This property of metal atoms, as you know, is determined by the fact that they have relatively large radii and a small number of electrons (mostly from $1$ to $3$ in the outer layer).

The only exceptions are $6$ metals: germanium, tin, and lead atoms on the outer layer have $4$ electrons, antimony and bismuth atoms have $5$, polonium atoms have $6$.

Metal atoms are characterized by low electronegativity values ​​(from $0.7$ to $1.9$) and exclusively reducing properties, i.e. ability to donate electrons.

You already know that in D.I. Mendeleev’s Periodic Table of Chemical Elements, metals are located below the boron-astatine diagonal, as well as above it, in secondary subgroups. In the periods and main subgroups, there are known patterns in the changes in the metallic, and therefore the reducing properties of the atoms of the elements.

Chemical elements located near the boron-astatine diagonal ($Be, Al, Ti, Ge, Nb, Sb$) have dual properties: in some of their compounds they behave like metals, in others they exhibit the properties of non-metals.

In secondary subgroups, the reducing properties of metals most often decrease with increasing atomic number.

This can be explained by the fact that the strength of the bond between the valence electrons and the nucleus of the atoms of these metals is influenced to a greater extent by the magnitude of the nuclear charge, and not by the radius of the atom. The nuclear charge increases significantly, and the attraction of electrons to the nucleus increases. In this case, although the atomic radius increases, it is not as significant as for the metals of the main subgroups.

Simple substances formed by chemical elements - metals, and complex metal-containing substances play a vital role in the mineral and organic “life” of the Earth. Suffice it to remember that atoms (ions) of metal elements are an integral part of compounds that determine metabolism in the body of humans and animals. For example, $76$ of elements were found in human blood, of which only $14$ are not metals. In the human body, some elements - metals (calcium, potassium, sodium, magnesium) are present in large quantities, i.e. are macroelements. And metals such as chromium, manganese, iron, cobalt, copper, zinc, molybdenum are present in small quantities, i.e. This microelements.

Features of the structure of metals of the main subgroups of groups I-III.

Alkali metals- these are metals of the main subgroup of group I. Their atoms at the outer energy level have one electron each. Alkali metals are strong reducing agents. Their reducing power and chemical activity increase with increasing atomic number of the element (i.e. from top to bottom in the Periodic Table). All of them have electronic conductivity. The strength of the bond between alkali metal atoms decreases with increasing atomic number of the element. Their melting and boiling points also decrease. Alkali metals interact with many simple substances - oxidizing agents. In reactions with water they form water-soluble bases (alkalis).

Alkaline earth elements are the elements of the main subgroup of group II. The atoms of these elements contain two electrons at the outer energy level. They are reducing agents and have an oxidation state of $+2$. In this main subgroup, general patterns in changes in physical and chemical properties are observed, associated with an increase in the size of atoms in the group from top to bottom, and the chemical bond between atoms also weakens. As the size of the ion increases, the acidic properties of oxides and hydroxides become weaker and the basic ones increase.

The main subgroup of group III consists of the elements boron, aluminum, gallium, indium and thallium. All elements are $p$-elements. At the outer energy level they have three $(s^2p^1)$ electrons, which explains the similarity of properties. Oxidation state $+3$. Within a group, as the nuclear charge increases, the metallic properties increase. Boron is a non-metallic element, while aluminum already has metallic properties. All elements form oxides and hydroxides.

Characteristics of transition elements - copper, zinc, chromium, iron according to their position in the Periodic Table of Chemical Elements of D. I. Mendeleev and the structural features of their atoms

Most metal elements are found in secondary groups of the Periodic Table.

In the fourth period, a fourth electron layer appears in the potassium and calcium atoms, and the $4s$ sublevel is filled, since it has lower energy than the $3d$ sublevel. $K, Ca are s$-elements included in the main subgroups. For atoms from $Sc$ to $Zn$, the $3d$ sublevel is filled with electrons.

Let's consider what forces act on an electron that is added to an atom as the nuclear charge increases. On the one hand, there is attraction by the atomic nucleus, which forces the electron to occupy the lowest free energy level. On the other hand, repulsion by already existing electrons. When there are $8$ electrons at the energy level (the $s-$ and $p-$orbitals are occupied), their overall repulsive effect is so strong that the next electron ends up in the higher $s-$ orbital instead of the energy level lower than the $d-$orbital. next level orbital. The electronic structure of the outer energy levels of potassium is $...3d^(0)4s^1$, and that of calcium is $...3d^(0)4s^2$.

The subsequent addition of one more electron to scandium leads to the beginning of the filling of the $3d$ orbital instead of even higher energy $4p$ orbitals. This turns out to be energetically more favorable. The filling of the $3d$ orbital ends in zinc, which has an electronic structure of $1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)3d^(10)4s^2$. It should be noted that the elements copper and chromium exhibit the phenomenon of electron “failure”. In a copper atom, the tenth $d$ electron moves to the third $3d$ sublevel.

The electronic formula of copper is $...3d^(10)4s^1$. A chromium atom in the fourth energy level ($s$-orbital) should have $2$ electrons. However, one of the two electrons moves to the third energy level, to an unfilled $d$-orbital, its electronic formula is $...3d^(5)4s^1$.

Thus, in contrast to the elements of the main subgroups, where the atomic orbitals of the outer level are gradually filled with electrons, the $d$-orbitals of the penultimate energy level are filled in the elements of the secondary subgroups. Hence the name: $d$-elements.

All simple substances formed by elements of subgroups of the Periodic Table are metals. Due to the greater number of atomic orbitals than those of the metal elements of the main subgroups, the atoms of the $d$ elements form a large number of chemical bonds with each other and therefore create a stronger crystal lattice. It is stronger both mechanically and in relation to heat. Therefore, metals of secondary subgroups are the strongest and most refractory among all metals.

It is known that if an atom has more than three valence electrons, then the element exhibits variable valence. This applies to most $d$ elements. Their maximum valence, like that of elements of the main subgroups, is equal to the group number (although there are exceptions). Elements with an equal number of valence electrons are included in the group under the same number $(Fe, Co, Ni)$.

For $d$-elements, the properties of their oxides and hydroxides change within one period when moving from left to right, i.e. with an increase in their valency, it proceeds from basic properties through amphoteric to acidic. For example, chromium has valencies $+2, +3, +6$; and its oxides: $CrO$ - basic, $Cr_(2)O_3$ - amphoteric, $CrO_3$ - acidic.

General characteristics of non-metals of the main subgroups of groups IV±VII in connection with their position in the Periodic Table of Chemical Elements of D. I. Mendeleev and the structural features of their atoms

Chemical elements - non-metals

The very first scientific classification of chemical elements was their division into metals and non-metals. This classification has not lost its significance to this day.

Nonmetals- these are chemical elements whose atoms are characterized by the ability to accept electrons before the completion of the outer layer due to the presence, as a rule, of four or more electrons on the outer electronic layer and the small radius of the atoms compared to metal atoms.

This definition leaves aside the elements of group VIII of the main subgroup - inert, or noble, gases, the atoms of which have a complete outer electron layer. The electronic configuration of the atoms of these elements is such that they cannot be classified as either metals or non-metals. They are those objects that divide elements into metals and non-metals, occupying a borderline position between them. Inert, or noble, gases (“nobility” is expressed in inertness) are sometimes classified as non-metals, but formally, according to their physical characteristics. These substances retain a gaseous state down to very low temperatures. Thus, helium He turns into a liquid state at $t°= -268.9 °C$.

The chemical inertness of these elements is relative. For xenon and krypton, compounds with fluorine and oxygen are known: $KrF_2, XeF_2, XeF_4$, etc. Undoubtedly, in the formation of these compounds, inert gases acted as reducing agents.

From the definition of nonmetals it follows that their atoms are characterized by high electronegativity values. It varies from $2$ to $4$. Nonmetals are elements of the main subgroups, mainly $p$-elements, with the exception of hydrogen, an s-element.

All non-metal elements (except hydrogen) occupy the upper right corner in the Periodic Table of Chemical Elements of D.I. Mendeleev, forming a triangle, the vertex of which is fluorine $F$, and the base is the diagonal $B - At$.

However, special attention should be paid to the dual position of hydrogen in the Periodic Table: in the main subgroups of groups I and VII. This is no coincidence. On the one hand, the hydrogen atom, like alkali metal atoms, has one electron on its outer (and only) electron layer (electronic configuration $1s^1$), which it is able to donate, exhibiting the properties of a reducing agent.

In most of its compounds, hydrogen, like alkali metals, exhibits an oxidation state of $+1$. But the loss of an electron by a hydrogen atom is more difficult than that of alkali metal atoms. On the other hand, the hydrogen atom, like the halogen atoms, lacks one electron before completing the outer electron layer, so the hydrogen atom can accept one electron, exhibiting the properties of an oxidizing agent and the oxidation state characteristic of a halogen - $1$ in hydrides (compounds with metals, similar to compounds metals with halogens - halides). But the addition of one electron to a hydrogen atom is more difficult than for halogens.

Properties of atoms of elements - nonmetals

Non-metal atoms have predominant oxidizing properties, i.e. ability to add electrons. This ability is characterized by the value of electronegativity, which naturally changes in periods and subgroups.

Fluorine is the strongest oxidizing agent; its atoms in chemical reactions are not capable of donating electrons, i.e. exhibit restorative properties.

Configuration of the outer electronic layer.

Other non-metals may exhibit reducing properties, although to a much weaker extent compared to metals; in periods and subgroups, their reducing ability changes in the opposite order compared to the oxidative ability.

Non-metal chemical elements only $16$! Quite a bit, considering that $114$ of elements are known. Two non-metal elements make up $76%$ of the mass of the earth's crust. These are oxygen ($49%$) and silicon ($27%$). The atmosphere contains $0.03%$ of the mass of oxygen in the earth's crust. Nonmetals make up $98.5%$ of the mass of plants, $97.6%$ of the mass of the human body. Nonmetals $C, H, O, N, S, P$ are organogens that form the most important organic substances of a living cell: proteins, fats, carbohydrates, nucleic acids. The composition of the air we breathe includes simple and complex substances, also formed by non-metal elements (oxygen $O_2$, nitrogen $N_2$, carbon dioxide $CO_2$, water vapor $H_2O$, etc.).

Hydrogen is the main element of the Universe. Many space objects (gas clouds, stars, including the Sun) consist of more than half hydrogen. On Earth, including the atmosphere, hydrosphere and lithosphere, it is only $0.88%$. But this is by mass, and the atomic mass of hydrogen is very small. Therefore, its small content is only apparent, and out of every $100$ atoms on Earth, $17$ are hydrogen atoms.

Chemistry tickets grade 9 with answers

Ticket No. 1

Periodic law and periodic system of chemical elements by D. I. Mendeleev. Patterns of changes in the properties of elements of small periods and main subgroups depending on their serial (atomic) number.

The periodic table has become one of the most important sources of information about chemical elements and the simple substances and compounds they form.

Dmitry Ivanovich Mendeleev created the Periodic Table while working on his textbook “Fundamentals of Chemistry,” achieving maximum consistency in the presentation of the material. The pattern of changes in the properties of the elements forming a system is called the Periodic Law.

According to the periodic law, formulated by Mendeleev in 1869, the properties of chemical elements are periodically dependent on their atomic masses. That is, with an increase in relative atomic mass, the properties of elements periodically repeat.*

Compare: the frequency of changing seasons over time.

This pattern is sometimes violated, for example, argon (an inert gas) exceeds the next potassium (alkali metal) in mass. This contradiction was explained in 1914 when studying the structure of the atom. The serial number of an element in the Periodic Table is not just a sequence, it has a physical meaning - it is equal to the charge of the nucleus of the atom. That's why

The modern formulation of the Periodic Law is:

The properties of chemical elements, as well as the substances formed by them, periodically depend on the charge of the atomic nucleus.

A period is a sequence of elements arranged in order of increasing charge on the nucleus of an atom, starting with an alkali metal and ending with an inert gas.

During the period, with an increase in the charge of the nucleus, the electronegativity of the element increases, the metallic (reducing) properties weaken and the non-metallic (oxidizing) properties of simple substances increase. Thus, the second period begins with the alkali metal lithium, followed by beryllium, which exhibits amphoteric properties, boron, a non-metal, etc. Finally, fluorine is a halogen and neon is an inert gas.

(The third period begins again with an alkali metal - this is periodicity)

Periods 1-3 are small (contain one row: 2 or 8 elements), periods 4-7 are large periods, consisting of 18 or more elements.

When compiling the periodic table, Mendeleev combined elements known at that time that had similarities into vertical columns. Groups are vertical columns of elements that, as a rule, have a valency in the higher oxide equal to the group number. The group is divided into two subgroups:

The main subgroups contain elements of small and large periods and form families with similar properties (alkali metals - I A, halogens - VII A, inert gases - VIII A).

(the chemical signs of the elements of the main subgroups in the periodic table are located under the letter “A” or, in very old tables where there are no letters A and B - under the element of the second period)

Side subgroups contain elements only of long periods; they are called transition metals.

(under the letter "B" or "B")

In the main subgroups, with increasing nuclear charge (atomic number), metallic (reducing) properties increase.

* more precisely, substances formed by elements, but this is often omitted when saying “properties of the elements”

Periodic law D.I. Mendeleev is a fundamental law that establishes a periodic change in the properties of chemical elements depending on the increase in the charges of the nuclei of their atoms. Opened by D.I. Mendeleev in March 1869 when comparing the properties of all elements known at that time and the values ​​of their atomic masses. Mendeleev first used the term “periodic law” in November 1870, and in October 1871 he gave the final formulation of the Periodic Law: “the properties of simple bodies, as well as the forms and properties of compounds of elements, and therefore the properties of the simple and complex bodies formed by them, stand in periodic dependence on their atomic weight."

Hund's rule: atomic orbitals belonging to the same sublevel are each filled first with one electron, and then they are filled with second electrons.

Hund's rule is also called the principle of maximum multiplicity, i.e. the maximum possible parallel direction of the spins of electrons of one energy sublevel.

A free atom can have no more than eight electrons at its highest energy level.

Electrons located at the highest energy level of an atom (in the outer electron layer) are called external; The number of outer electrons in an atom of any element is never more than eight. For many elements, it is the number of external electrons (with filled internal sublevels) that largely determines their chemical properties. For other electrons whose atoms have an unfilled internal sublevel, for example 3 d- sublevel of atoms of elements such as Sc, Ti, Cr, Mn, etc., chemical properties depend on the number of both internal and external electrons. All these electrons are called valence; in abbreviated electronic formulas of atoms they are written after the symbol of the atomic skeleton, that is, after the expression in square brackets.

2.3. Periodic Law and Periodic Table of Elements

At the beginning of the 20th century, with the discovery of the structure of the atom, it was established that the periodicity of changes in the properties of elements is determined not atomic weight, but nuclear charge, equal to the atomic number and the number of electrons, the distribution of which over the electron shells of an element’s atom determines its chemical properties.

The modern formulation of the Periodic Law reads:

The further development of the periodic system is associated with filling the empty cells of the table, into which more and more new elements were placed: noble gases, natural and artificially obtained radioactive elements. In 2010, with the synthesis of element 117, the seventh period of the periodic table was completed. However, the problem of the lower limit of the periodic table remains one of the most important in modern theoretical chemistry

The graphical (tabular) expression of the periodic law is the periodic system of elements developed by Mendeleev .

More common than others are 3 forms of the periodic table: “short” (short-period), “long” (long-period), “extra-long”.

In the “super-long” version, each period occupies exactly one line. In the “long” version, lanthanides and actinides are removed from the general table, making it more compact. In the “short” form of recording, in addition to this, the fourth and subsequent periods occupy 2 lines each.

Elements arranged in ascending order of Z (H, He, Li, Be...) form seven periods.

In periods the properties of elements naturally change during the transition from alkali metals to noble gases