What is the physical meaning of the group number? Periodic table of elements

The concept of elements as primary substances dates back to ancient times and, gradually changing and becoming more precise, has reached our time. The founders of scientific views on chemical elements are R. Boyle (7th century), M.V. Lomonosov (18th century) and Dalton (19th century).
By the beginning of the 19th century. About 30 elements were known, by the middle of the 19th century - about 60. As the number of elements accumulated, the task of systematizing them arose. Such attempts before D.I. Mendeleev was no less than fifty; The systematization was based on: atomic weight (now called atomic mass), chemical equivalent, and valence. Approaching the classification of chemical elements metaphysically, trying to systematize only the elements known at that time, none of D.I. Mendeleev’s predecessors could discover the universal interconnection of elements or create a single harmonious system reflecting the law of development of matter. This important problem for science was brilliantly solved in 1869 by the great Russian scientist D.I. Mendeleev, who discovered the periodic law.
Mendeleev’s systematization was based on: a) atomic weight and b) chemical similarity between elements. The most striking expression of the similarity of the properties of elements is their identical highest valence. Both atomic weight (atomic mass) and the highest valence of an element are quantitative, numerical constants convenient for systematization.
Having arranged all 63 elements known at that time in a row in order of increasing atomic masses, Mendeleev noticed the periodic repeatability of the properties of elements at unequal intervals. As a result, Mendeleev created the first version of the periodic table.
The regular nature of the change in the atomic masses of elements along the verticals and horizontals of the table, as well as the empty spaces formed in it, allowed Mendeleev to boldly predict the presence in nature of a number of elements that were not yet known to science at that time and even outline their atomic masses and basic properties based on the expected position elements in the table. This could only be done on the basis of a system that objectively reflects the law of development of matter. The essence of the periodic law D.I. Mendeleev formulated in 1869: “The properties of simple bodies, as well as the forms and properties of compounds of elements are periodically dependent on the magnitude of the atomic weights (mass) of the elements.”

Periodic table of elements.
In 1871, D.I. Mendeleev gives the second version of the periodic table (the so-called short form of the table), in which he identifies various degrees of relationship between elements. This version of the system made it possible for Mendeleev to predict the existence of 12 elements and describe the properties of three of them with very high accuracy. In the period from 1875 to 1886. these three elements were discovered and a complete coincidence of their properties with those predicted by the great Russian scientist was revealed. These elements received the following names: scandium, gallium, germanium. After this, the periodic law received universal recognition as an objective law of nature and is now the foundation of chemistry, physics and other natural sciences.

The periodic table of chemical elements is a graphic expression of the periodic law. It is known that a number of laws, in addition to verbal formulations, can be depicted graphically and expressed in mathematical formulas. This is also the periodic law; only the mathematical laws inherent in it, which will be discussed below, are not yet united by a general formula. Knowledge of the periodic table makes it easier to study general chemistry.
The design of the modern periodic system, in principle, differs little from the version of 1871. The symbols of the elements in the periodic system are arranged in vertical and horizontal columns. This leads to the unification of elements into groups, subgroups, periods. Each element occupies a specific cell in the table. Vertical graphs are groups (and subgroups), horizontal graphs are periods (and series).

By group is a collection of elements with the same oxygen valency. This highest valence is determined by the group number. Since the sum of the highest valences for oxygen and hydrogen for non-metal elements is eight, it is easy to determine the formula of the highest hydrogen compound by the group number. So, for phosphorus - an element of the fifth group - the highest valency for oxygen is five, the formula of the highest oxide is P2O5, and the formula of the compound with hydrogen is PH3. For sulfur, an element of the sixth group, the formula of the highest oxide is SO3, and the formula of the highest compound with hydrogen is H2S.
Some elements have a higher valency that is not equal to their group number. Such exceptions are copper Cu, silver Ag, gold Au. They are in the first group, but their valencies vary from one to three. For example, there are compounds: CuO; AgO; Cu2O3; Au2O3. Oxygen is placed in the sixth group, although its compounds with a valency higher than two are almost never found. Fluorine P, an element of group VII, is monovalent in its most important compounds; Bromine Br, an element of group VII, is maximally pentavalent. There are especially many exceptions in group VIII. There are only two elements in it: ruthenium Ru and osmium Os exhibit a valence of eight; their higher oxides have the formulas RuO4 and OsO4. The valence of the remaining elements of group VIII is much lower.
Initially, Mendeleev's periodic system consisted of eight groups. At the end of the 19th century. Inert elements predicted by the Russian scientist N.A. Morozov were discovered, and the periodic table was replenished with a ninth group - number zero. Now many scientists consider it necessary to return to dividing all elements into 8 groups again. This makes the system more harmonious; From the perspective of the octet (eight) groups, some rules and laws become clearer.

Group elements are distributed according to subgroups. A subgroup combines elements of a given group that are more similar in their chemical properties. This similarity depends on the analogy in the structure of the electronic shells of the atoms of the elements. In the periodic table, the symbols of the elements of each subgroup are arranged strictly vertically.
The first seven groups have one main and one secondary subgroup; in the eighth group there is one main subgroup, “inert” elements, and three secondary ones. The name of each subgroup is usually given by the name of the top element, for example: lithium subgroup (Li-Na-K-Rb-Cs-Fr), chromium subgroup (Cr-Mo-W). While elements of the same subgroup are chemical analogues, elements of different subgroups of the same group sometimes differ very sharply in their properties. The common property for elements of the main and secondary subgroups of the same group is basically only their identical highest oxygen valency. Thus, manganese Mn and chlorine C1, located in different subgroups of group VII, chemically have almost nothing in common: manganese is a metal, chlorine is a typical non-metal. However, the formulas of their higher oxides and the corresponding hydroxides are similar: Mn2O7 - Cl2O7; НМnО4 - НС1О4.
The periodic table has two horizontal rows of 14 elements located outside the groups. They are usually placed at the bottom of the table. One of these series consists of elements called lanthanides (literally: like lanthanum), the other series consists of elements called actinides (like actinium). The actinide symbols are located below the lanthanide symbols. This arrangement reveals 14 shorter subgroups, consisting of 2 elements each: these are the second secondary, or lanthanide-actinide subgroups.
Based on all that has been said, they distinguish: a) main subgroups, b) secondary subgroups and c) second secondary (lanthanide-actinide) subgroups.

It should be taken into account that some main subgroups also differ from each other in the structure of the atoms of their elements. Based on this, all subgroups of the periodic table can be divided into 4 categories.
I. Main subgroups of groups I and II (subgroups of lithium and beryllium).
II. Six main subgroups III - IV - V - VI - VII - VIII groups (subgroups of boron, carbon, nitrogen, oxygen, fluorine and neon).
III. Ten side subgroups (one in groups I - VII and three in group VIII). jfc,
IV. Fourteen lanthanide-actinide subgroups.
The numbers of subgroups of these 4 categories form an arithmetic progression: 2-6-10-14.
It should be noted that the top element of any main subgroup is in period 2; the top element of any side element - in the 4th period; the top element of any lanthanide-actinide subgroup - in the 6th period. Thus, with each new even period of the periodic table, new categories of subgroups appear.
Each element, in addition to being in one or another group and subgroup, is also located in one of the seven periods.
A period is a sequence of elements during which their properties change in order of gradual intensification from typically metallic to typically nonmetallic (metalloid). Each period ends with an inert element. As the metallic properties of the elements weaken, non-metallic properties begin to appear and gradually increase; in the middle of the periods there are usually elements that combine, to one degree or another, both metallic and non-metallic properties. These elements are often called amphoteric.

Composition of periods.
The periods are not uniform in the number of elements included in them. The first three are called small, the remaining four are called large. In Fig. Figure 8 shows the composition of the periods. The number of elements in any period is expressed by the formula 2n2 where n is an integer. Periods 2 and 3 contain 8 elements each; in 4 and 5 - 18 elements each; in 6-32 elements; in 7, which is not yet finished, there are 18 elements, although theoretically there should also be 32 elements.
Original 1st period. It contains only two elements: hydrogen H and helium He. The transition of properties from metallic to non-metallic occurs here in one typically amphoteric element - hydrogen. The latter, in terms of its inherent metallic properties, heads the subgroup of alkali metals, and in terms of its inherent non-metallic properties, it heads the subgroup of halogens. Hydrogen is therefore often placed in the periodic table twice - in groups 1 and 7.

The different quantitative composition of the periods leads to an important consequence: neighboring elements of small periods, for example, carbon C and nitrogen N, differ relatively sharply from each other in their properties: neighboring elements of large periods, for example, lead Pb and bismuth Bi, are much closer in properties to each other friend, since the change in the nature of the elements over long periods occurs in small jumps. In certain areas of long periods, there is even such a slow decline in metallicity that nearby elements turn out to be very similar in their chemical properties. This is, for example, the triad of elements of the fourth period: iron Fe - cobalt Co - nickel Ni, which is often called the “iron family”. The horizontal similarity (horizontal analogy) here even overlaps the vertical similarity (vertical analogy); Thus, the elements of the iron subgroup - iron, ruthenium, osmium - are less chemically similar to each other than the elements of the “iron family”.
The most striking example of a horizontal analogy is the lanthanides. All of them are chemically similar to each other and to lanthanum La. In nature, they are found in groups, difficult to separate, the typical highest valence of most of them is 3. Lanthanides have a special internal periodicity: every eighth of them, in order of arrangement, repeats to some extent the properties and valence states of the first, i.e. the one from which the countdown begins. Thus, terbium Tb is similar to cerium Ce; lutetium Lu - to gadolinium Gd.
Actinides are similar to lanthanides, but their horizontal analogy is much less pronounced. The highest valency of some actinides (for example, uranium U) reaches six. The fundamentally possible internal periodicity among them has not yet been confirmed.

Arrangement of elements in the periodic table. Moseley's Law.

D.I. Mendeleev arranged the elements in a certain sequence, sometimes called the “Mendeleev series”. In general, this sequence (numbering) is associated with an increase in the atomic masses of the elements. However, there are exceptions. Sometimes the logical course of changes in valency is in conflict with the course of changes in atomic masses . In such cases, necessity required giving preference to one of these two principles of systematization. D. I. Mendeleev in some cases violated the principle of the arrangement of elements with increasing atomic masses and relied on the chemical analogy between the elements. If Mendeleev had placed nickel Ni in front of cobalt Co, iodine I before tellurium Te, then these elements would fall into subgroups and groups that do not correspond to their properties and their highest valency.
In 1913, the English scientist G. Moseley, studying the spectra of X-rays for various elements, noticed a pattern connecting the numbers of elements in Mendeleev’s periodic table with the wavelength of these rays resulting from the irradiation of certain elements by cathode clouds. It turned out that the square roots of the reciprocal wavelengths of these rays are linearly related to the serial numbers of the corresponding elements. G. Moseley's law made it possible to verify the correctness of the “Mendeleev series” and confirmed its impeccability.
Let, for example, know the values for elements No. 20 and No. 30, the numbers of which in the system do not cause us any doubt. These values are related to the indicated numbers by a linear relationship. To check, for example, that the number assigned to cobalt (27) is correct, and judging by the atomic mass, this number should have been nickel, it is irradiated with cathode rays: as a result, X-rays are released from cobalt. By decomposing them on suitable diffraction gratings (crystals), we obtain the spectrum of these rays and, choosing the clearest of the spectral lines, we measure the wavelength () of the ray corresponding to this line; then we plot the value on the ordinate. From the resulting point A, draw a straight line parallel to the x-axis until it intersects with the previously identified straight line. From intersection point B, we lower the perpendicular to the x-axis: it will accurately indicate to us the cobalt number, equal to 27. Thus, the periodic system of elements of D. I. Mendeleev - the fruit of the scientist’s logical conclusions - received experimental confirmation.

Modern formulation of the periodic law. The physical meaning of the element's serial number.

After the work of G. Moseley, the atomic mass of an element gradually began to give up its primacy role to a new, not yet clear in its internal (physical) meaning, but a clearer constant - the ordinal or, as they now call it, the atomic number of the element. The physical meaning of this constant was revealed in 1920 by the work of the English scientist D. Chadwick. D. Chadwick experimentally established that the atomic number of an element is numerically equal to the positive charge Z of the nucleus of an atom of this element, i.e., the number of protons in the nucleus. It turned out that D.I. Mendeleev, without suspecting it, arranged the elements in a sequence that exactly corresponded to the increase in the charge of the nuclei of their atoms.
By this time it was also established that atoms of the same element can differ from each other in their mass; such atoms are called isotopes. An example would be atoms: and . In the periodic table, isotopes of the same element occupy one cell. In connection with the discovery of isotopes, the concept of a chemical element was clarified. Currently, a chemical element is a type of atoms that have the same nuclear charge - the same number of protons in the nucleus. The formulation of the periodic law was also clarified. The modern formulation of the law states: the properties of elements and their compounds are periodically dependent on the size and charge of the nuclei of their atoms.
Other characteristics of elements related to the structure of the outer electronic layers of atoms, atomic volumes, ionization energy and other properties also change periodically.

Periodic system and structure of electronic shells of atoms of elements.

Later it was found that not only the serial number of an element has a deep physical meaning, but also other concepts previously discussed also gradually acquired a physical meaning. For example, the group number, indicating the highest valence of an element, thereby reveals the maximum number of electrons in an atom of a particular element that can participate in the formation of a chemical bond.
The period number, in turn, turned out to be related to the number of energy levels present in the electron shell of an atom of an element of a given period.
Thus, for example, the “coordinates” of tin Sn (serial number 50, period 5, main subgroup of group IV) mean that there are 50 electrons in a tin atom, they are distributed over 5 energy levels, only 4 electrons are valence.
The physical meaning of finding elements in subgroups of various categories is extremely important. It turns out that for elements located in category I subgroups, the next (last) electron is located on the s-sublevel of the outer level. These elements belong to the electronic family. For atoms of elements located in subgroups of category II, the next electron is located at the p-sublevel of the outer level. These are elements of the “p” electronic family. Thus, the next 50th electron in tin atoms is located on the p-sublevel of the external, i.e., 5th energy level.
For atoms of elements of category III subgroups, the next electron is located on the d-sublevel, but already before the outer level, these are elements of the “d” electronic family. In lanthanide and actinide atoms, the next electron is located on the f-sublevel, before the outer level. These are elements of the “f” electronic family.
It is no coincidence, therefore, that the numbers of subgroups of these 4 categories noted above, that is, 2-6-10-14, coincide with the maximum numbers of electrons in the s-p-d-f sublevels.
But it turns out that it is possible to solve the question of the order of filling the electron shell and derive the electronic formula for an atom of any element on the basis of the periodic system, which with sufficient clarity indicates the level and sublevel of each successive electron. The periodic system also indicates the placement of elements one after another into periods, groups, subgroups and the distribution of their electrons among levels and sublevels, because each element has its own, characterizing its last electron. As an example, let's look at compiling an electronic formula for an atom of the element zirconium (Zr). The periodic system gives indicators and “coordinates” of this element: serial number 40, period 5, group IV, secondary subgroup. First conclusions: a) there are 40 electrons in all, b) these 40 electrons are distributed at five energy levels; c) out of 40 electrons only 4 are valence, d) the next 40th electron entered the d-sublevel before the outer, i.e., fourth energy level.Similar conclusions can be drawn about each of the 39 elements preceding zirconium, only the indicators and coordinates will be different each time.
Therefore, the methodological technique for compiling electronic formulas of elements based on the periodic system is that we sequentially consider the electronic shell of each element along the way to a given one, identifying by its “coordinates” where its next electron went in the shell.
The first two elements of the first period, hydrogen H and helium He, belong to the s-family. Two of their electrons enter the s-sublevel of the first level. We write down: The first period ends here, the first energy level also. The next two elements in order of the second period - lithium Li and beryllium Be are in the main subgroups of groups I and II. These are also s-elements. Their next electrons will be located on the s sublevel of the 2nd level. We write down 6 elements of the 2nd period follow in a row: boron B, carbon C, nitrogen N, oxygen O, fluorine F and neon Ne. According to the location of these elements in the main subgroups of the III - Vl groups, their next electrons, among the six, will be located on the p-sublevel of the 2nd level. We write down: The inert element neon ends the second period, the second energy level is also completed. This is followed by two elements of the third period of the main subgroups of groups I and II: sodium Na and magnesium Mg. These are s-elements and their next electrons are located on the s-sublevel of the 3rd level. Then there are six elements of the 3rd period: aluminum Al, silicon Si, phosphorus P, sulfur S, chlorine C1, argon Ar. According to the location of these elements in the main subgroups of groups III - UI, their next electrons, among the six, will be located on the p-sublevel of the 3rd level - The inert element argon has completed the 3rd period, but the 3rd energy level has not yet been completed, as long as there are no electrons on its third possible d-sublevel.
This is followed by 2 elements of the 4th period of the main subgroups of groups I and II: potassium K and calcium Ca. These are s-elements again. Their next electrons will be at the s-sublevel, but already at the 4th level. It is energetically more favorable for these next electrons to start filling the 4th level, which is more distant from the nucleus, than to fill the 3d sublevel. We write down: The following ten elements of the 4th period from No. 21 scandium Sc to No. 30 zinc Zn are in secondary subgroups III - V - VI - VII - VIII - I - II groups. Since they are all d-elements, their next electrons are located on the d-sublevel before the outer level, i.e., the third from the nucleus. We write down:
The following six elements of the 4th period: gallium Ga, germanium Ge, arsenic As, selenium Se, bromine Br, krypton Kr - are in the main subgroups of groups III - VIIJ. Their next 6 electrons are located on the p-sublevel of the outer, i.e., 4th level: 3b elements were considered; the fourth period is completed by the inert element krypton; The 3rd energy level is also completed. However, at level 4, only two sublevels are completely filled: s and p (out of 4 possible).
This is followed by 2 elements of the 5th period of the main subgroups of groups I and II: No. 37 rubidium Rb and No. 38 strontium Sr. These are elements of the s-family, and their next electrons are located on the s-sublevel of the 5th level: The last 2 elements - No. 39 yttrium YU No. 40 zirconium Zr - are already in secondary subgroups, i.e. they belong to the d-family. Their next two electrons will go to the d-sublevel, before the outer one, i.e. 4th level Summing up all the records sequentially, we compose the electronic formula for the zirconium atom No. 40. The derived electronic formula for the zirconium atom can be slightly modified by arranging the sublevels in the order of numbering their levels:

The derived formula can, of course, be simplified into the distribution of electrons only among energy levels: Zr – 2|8| 18 |8 + 2| 2 (the arrow indicates the entry point of the next electron; valence electrons are underlined). The physical meaning of the category of subgroups lies not only in the difference in the place where the next electron enters the shell of the atom, but also in the levels at which the valence electrons are located. From a comparison of simplified electronic formulas, for example, chlorine (3rd period, main subgroup of group VII), zirconium (5th period, secondary subgroup of group IV) and uranium (7th period, lanthanide-actinide subgroup)
№17, С1-2|8|7
No. 40, Zr - 2|8|18|8+ 2| 2
No. 92, U - 2|8|18 | 32 |18 + 3|8 + 1|2
It can be seen that for elements of any main subgroup, only electrons of the outer level (s and p) can be valence. For elements of side subgroups, valence electrons can be the electrons of the outer and partially pre-outer level (s and d). In lanthanides and especially actinides, valence electrons can be located at three levels: external, pre-external and pre-external. Typically, the total number of valence electrons is equal to the group number.

Element properties. Ionization energy. Electron affinity energy.

A comparative examination of the properties of elements is carried out in three possible directions of the periodic system: a) horizontal (by period), b) vertical (by subgroup), c) diagonal. To simplify our reasoning, we will exclude the 1st period, the unfinished 7th period, as well as the entire VIII group. The main parallelogram of the system will remain, in the upper left corner of which there will be lithium Li (No. 3), in the lower left - cesium Cs (No. 55). In the upper right - fluorine F (No. 9), in the lower right - astatine At (No. 85).
directions. In the horizontal direction from left to right, the volumes of atoms gradually decrease; occurs, this is as a result of the influence of an increase in the charge of the nucleus on the electron shell. In the vertical direction from top to bottom, as a result of increasing the number of levels, the volumes of atoms gradually increase; along the diagonal direction - much less clearly defined and shorter - remain close. These are general patterns, to which, as always, there are exceptions.
In the main subgroups, as the volume of atoms increases, i.e., from top to bottom, the detachment of external electrons becomes easier and the addition of new electrons to the atoms becomes more difficult. The donation of electrons characterizes the so-called reducing power of elements, especially typical of metals. The addition of electrons characterizes the oxidizing ability, which is typical for non-metals. Consequently, from top to bottom in the main subgroups, the reducing ability of atoms of elements increases; The metallic properties of simple bodies corresponding to these elements also increase. The oxidative capacity decreases.
From left to right across periods, the pattern of changes is the opposite: the reducing ability of elemental atoms decreases, while the oxidative ability increases; the non-metallic properties of simple bodies corresponding to these elements increase.
Along the diagonal direction, the properties of the elements remain more or less close. Let's look at this direction using an example: beryllium-aluminum
From beryllium Be to aluminum Al you can go directly along the diagonal Be → A1, or through boron B, that is, along two legs Be → B and B → A1. The strengthening of non-metallic properties from beryllium to boron and their weakening from boron to aluminum explains why the elements beryllium and aluminum, located on the diagonal, have some analogy in properties, although they are not in the same subgroup of the periodic table.
Thus, there is a close connection between the periodic table, the structure of atoms of elements and their chemical properties.
The properties of an atom of any element - giving up an electron and turning into a positively charged ion - are quantified by the expenditure of energy, called ionization energy I*. It is expressed in kcal/g-atom or hj/g-atom.

The lower this energy, the stronger the element’s atom exhibits reducing properties, the more metallic the element; The greater this energy, the weaker the metallic properties, the stronger the non-metallic properties of the element. The property of an atom of any element to accept an electron and transform into a negatively charged ion is assessed by the amount of energy released, called electron affinity E; it is also expressed in kcal/g-atom or kJ/g-atom.

Electron affinity can be a measure of an element's ability to exhibit nonmetallic properties. The greater this energy, the more non-metallic the element, and, conversely, the less energy, the more metallic the element.
Often, to characterize the properties of elements, a quantity called electronegativity.
It: is the arithmetic sum of the ionization energy and the electron affinity energy

The constant is a measure of the non-metallicity of elements. The larger it is, the stronger the element exhibits non-metallic properties.
It should be borne in mind that all elements are essentially dual in nature. The division of elements into metals and non-metals is to a certain extent arbitrary, since in nature there are no sharp edges. As an element's metallic properties increase, its non-metallic properties weaken and vice versa. The most “metallic” of the elements - francium Fr - can be considered the least non-metallic, the most “non-metallic” - fluorine F - can be considered the least metallic.
Summing up the values of the calculated energies - ionization energy and electron affinity energy - we obtain: for cesium the value is 90 kcal/g-a., for lithium 128 kcal/g-a., for fluorine = 510 kcal/g-a. (the value is also expressed in kJ/g-a.). These are absolute electronegativity values. To simplify, we use relative electronegativity values, taking the electronegativity of lithium (128) as unity. Then for fluorine (F) we get:

For cesium (Cs), the relative electronegativity will be equal to
On the graph of changes in electronegativity of elements of the main subgroups
I-VII groups. The electronegativities of the elements of the main subgroups of groups I-VII are compared. The given data indicate the true position of hydrogen in the 1st period; unequal increase in the metallicity of elements, from top to bottom in various subgroups; some similarity of elements: hydrogen - phosphorus - tellurium (= 2.1), beryllium and aluminum (= 1.5) and a number of other elements. As can be seen from the above comparisons, using electronegativity values, it is possible to approximately compare elements of even different subgroups and different periods with each other.

Graph of changes in the electronegativity of elements of the main subgroups of groups I-VII.

The periodic law and the periodic system of elements have enormous philosophical, scientific and methodological significance. They are: a means of understanding the world around us. The periodic law reveals and reflects the dialectical-materialistic essence of nature. The periodic law and the periodic system of elements convincingly prove the unity and materiality of the world around us. They are the best confirmation of the validity of the main features of the Marxist dialectical method of cognition: a) the interconnection and interdependence of objects and phenomena, b) continuity of movement and development, c) the transition of quantitative changes into qualitative ones, d) struggle and unity of opposites.
The enormous scientific significance of the periodic law lies in the fact that it helps creative discoveries in the field of chemical, physical, mineralogical, geological, technical and other sciences. Before the discovery of the periodic law, chemistry was an accumulation of scattered factual information devoid of internal connection; Now all this has been brought into a single harmonious system. Many discoveries in the field of chemistry and physics were made on the basis of the periodic law and the periodic table of elements. The periodic law opened the way to knowledge of the internal structure of the atom and its nucleus. It is enriched with ever new discoveries and is confirmed as an unshakable, objective law of nature. The great methodological and methodological significance of the periodic law and the periodic system of elements lies in the fact that when studying chemistry, they provide the opportunity to develop a dialectical-materialistic worldview in the student and facilitate the acquisition of a chemistry course: The study of chemistry should not be based on memorizing the properties of individual elements and their compounds, but judge the properties of simple and complex substances based on the patterns expressed by the periodic law and the periodic system of elements.

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

Modern wording:

“the properties of chemical elements (i.e., the properties and form of the compounds they form) are periodically dependent on the charge of the nucleus of the atoms of the chemical elements.”

Physical meaning of chemical periodicity

Periodic changes in the properties of chemical elements are caused by the correct repetition of the electronic configuration of the outer energy level (valence electrons) of their atoms with an increase in the charge of the nucleus.

A graphical 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 principal quantum number of valence electrons.

The period number indicates the number of energy levels in an atom of an element.

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 give up electrons, and non-metallic properties to gain electrons due to the desire of atoms to acquire a stable configuration with filled sublevels. Filling outer s- sublevel indicates the metallic properties of the atom, and the formation of the outer p- sublevel - on non-metallic properties. 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 large periods, the transition of properties from an active metal to a noble gas occurs more smoothly than in short periods, because formation of internal ( n - 1) d - sublevel while maintaining the external ns 2 - layer. Large periods consist of even and odd series.

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 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.

Side 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 table are divided into: s- elements (elements of the main subgroup Groups I and II), p - elements (elements of the main subgroups III - VII groups), d - elements (elements of side subgroups), f- elements (lanthanides, actinides).

In the main subgroups, from top to bottom, metallic properties increase, and non-metallic properties weaken. 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).

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

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

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

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

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

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

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

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

Model the construction of the Periodic Table of D.I. Mendeleev. To do this, prepare 20 cards measuring 6 x 10 cm for elements with serial numbers from 1st to 20th. On each card, indicate the following information about the element: chemical symbol, name, relative atomic mass, formula of higher oxide, hydroxide (indicate their nature in parentheses - basic, acidic or amphoteric), formula of volatile hydrogen compound (for non-metals).

Shuffle the cards and then arrange them in a row in order of increasing relative atomic masses of the elements. Place similar elements from 1st to 18th under each other: hydrogen above lithium and potassium under sodium, respectively, calcium under magnesium, helium under neon. Formulate the pattern you have identified in the form of a law. Note the discrepancy between the relative atomic masses of argon and potassium and their location in terms of the common properties of the elements. Explain the reason for this phenomenon.

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

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

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

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

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

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

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

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

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

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

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

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

table 2
Predicted and experimentally discovered properties of germanium

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

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

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

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

New words and concepts

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

Tasks for independent work

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

Option 1

A1. What is the physical meaning of the group number of D.I. Mendeleev’s table?

2.This is the charge of the nucleus of an atom

4.This is the number of neutrons in the nucleus

A2. What is the number of energy levels?

1. Serial number

2. Period number

3. Group number

4. Number of electrons

A3.

2. This is the number of energy levels in an atom

3. This is the number of electrons in an atom

A4. Indicate the number of electrons in the outer energy level in the phosphorus atom:

1. 7 electrons

2. 5 electrons

3. 2 electrons

4. 3 electrons

A5. In which row are the formulas of hydrides located?

1. H 2 O, CO, C 2 H 2 , LiH

2.NaH, CH 4 , H 2 O,CaH 2

3. H 2 O,C 2 H 2 , LiH, Li 2 O

4. NO, N 2 O 3 , N 2 O 5 , N 2 O

A 6. In which compound is the oxidation state of nitrogen equal to +1?

1. N 2 O 3

2. NO

3. N 2 O 5

4. N 2 O

A7. Which compound corresponds to manganese (II) oxide:

1. MnO 2

2. Mn 2 O 7

3. MnCl 2

4. MnO

A8. Which row contains only simple substances?

1. Oxygen and ozone

2. Sulfur and water

3. Carbon and bronze

4. Sugar and salt

A9. Identify an element if its atom has 44 electrons:

1. cobalt

2. tin

3. ruthenium

4. niobium

A10. What has an atomic crystal lattice?

1. iodine

2. germanium

3. ozone

4. white phosphorus

IN 1. Match

Number of electrons in the outer energy level of an atom

Chemical element symbol

A. 3

B. 1

AT 6

G. 4

1) S 6) C

2) Fr 7) He

3) Mg 8) Ga

4) Al 9) Te

5) Si 10) K

AT 2. Match

Substance name

Substance formula

A. Oxidesulfur(VI)

B. Sodium hydride

B. Sodium hydroxide

G. Iron(II) chloride

1)SO 2

2) FeCl 2

3) FeCl 3

4) NaH

5) SO 3

6) NaOH

Option 2

A1. What is the physical meaning of the period number of D.I. Mendeleev’s table?

1.This is the number of energy levels in an atom

2.This is the charge of the nucleus of an atom

3. This is the number of electrons in the outer energy level of an atom

4.This is the number of neutrons in the nucleus

A2. What is the number of electrons in an atom?

1. Serial number

2. Period number

3. Group number

4. Number of neutrons

A3. What is the physical meaning of the atomic number of a chemical element?

1. This is the number of neutrons in the nucleus

2. This is the charge of the atomic nucleus

3. This is the number of energy levels in an atom

4. This is the number of electrons in the outer energy level of an atom

A4. Indicate the number of electrons in the outer energy level in a silicon atom:

1. 14 electrons

2. 4 electrons

3. 2 electrons

4. 3 electrons

A5. In what row are the oxide formulas located?

1. H 2 O, CO, CABOUT 2 , LiABOUTH

2.NaH, CH 4 , H 2 O,CaH 2

3. H 2 O,C 2 H 2 , LiH, Li 2 O

4. NO, N 2 O 3 , N 2 O 5 , N 2 O

A 6. In which compound does the oxidation state of chlorine equal to -1?

1. Cl 2 O 7

2. HClO

3. HCl

4. Cl 2 O 3

A7. Which compound corresponds to nitric oxide (III):

1. N 2 O

2. N 2 O 3

3. NO

4. H 3 N

A8. In which row are simple and complex substances located?

1. Diamond and ozone

2. Gold and carbon dioxide

3. Water and sulfuric acid

4. Sugar and salt

A9. Identify an element if its atom has 56 protons:

1. iron

2. tin

3. barium

4. manganese

A10. What has a molecular crystal lattice?

diamond

silicon

rhinestone

boron

IN 1. Match

Number of energy levels in an atom

Chemical element symbol

A. 5

B. 7

IN. 3

G. 2

1) S 6) C

2) Fr 7) He

3) Mg 8) Ga

4) B 9) Te

5) Sn 10) Rf

AT 2. Match

Substance name

Substance formula

A. Carbon hydride (IV)

B. Calcium oxide

B. Calcium nitride

G. Calcium hydroxide

1) H 3 N

2) Ca(OH) 2

3) KOH

4) CaO

5)CH 4

6) Ca 3 N 2