Physical foundations of the periodic table of chemical elements. Structure of the periodic table of Mendeleev. What is a period

Periodic law D.I. Mendeleev:Properties of simple bodies, as well as shapes and properties of compoundsdifferences of elements are periodically dependent onthe values of the atomic weights of elements. (The properties of elements are periodically dependent on the charge of the atoms of their nuclei).

Periodic table of elements. Series of elements within which properties change sequentially, such as the series of eight elements from lithium to neon or from sodium to argon, Mendeleev called periods. If we write these two periods one below the other so that sodium is under lithium and argon is under neon, we get the following arrangement of elements:

With this arrangement, the vertical columns contain elements that are similar in their properties and have the same valency, for example, lithium and sodium, beryllium and magnesium, etc.

Having divided all the elements into periods and placing one period under another so that elements similar in properties and type of compounds formed were located under each other, Mendeleev compiled a table that he called the periodic system of elements by groups and series.

The meaning of the periodic systemWe. The periodic table of elements had a great influence on the subsequent development of chemistry. Not only was it the first natural classification of chemical elements, showing that they form a harmonious system and are in close connection with each other, but it was also a powerful tool for further research.

7. Periodic changes in the properties of chemical elements. Atomic and ionic radii. Ionization energy. Electron affinity. Electronegativity.

The dependence of atomic radii on the charge of the nucleus of an atom Z is periodic. Within one period, as Z increases, there is a tendency for the size of the atom to decrease, which is especially clearly observed in short periods

With the beginning of the construction of a new electronic layer, more distant from the nucleus, i.e., during the transition to the next period, atomic radii increase (compare, for example, the radii of fluorine and sodium atoms). As a result, within a subgroup, with increasing nuclear charge, the sizes of atoms increase.

The loss of electron atoms leads to a decrease in its effective size, and the addition of excess electrons leads to an increase. Therefore, the radius of a positively charged ion (cation) is always smaller, and the radius of a negatively charged non (anion) is always greater than the radius of the corresponding electrically neutral atom.

Within one subgroup, the radii of ions of the same charge increase with increasing nuclear charge. This pattern is explained by an increase in the number of electronic layers and the growing distance of outer electrons from the nucleus.

The most characteristic chemical property of metals is the ability of their atoms to easily give up external electrons and transform into positively charged ions, while non-metals, on the contrary, are characterized by the ability to add electrons to form negative ions. To remove an electron from an atom and transform the latter into a positive ion, it is necessary to expend some energy, called ionization energy.

Ionization energy can be determined by bombarding atoms with electrons accelerated in an electric field. The lowest field voltage at which the electron speed becomes sufficient to ionize atoms is called the ionization potential of the atoms of a given element and is expressed in volts. With the expenditure of sufficient energy, two, three or more electrons can be removed from an atom. Therefore, they speak of the first ionization potential (the energy of the removal of the first electron from the atom) and the second ionization potential (the energy of the removal of the second electron)

As noted above, atoms can not only donate, but also gain electrons. The energy released when an electron is added to a free atom is called the atom's electron affinity. Electron affinity, like ionization energy, is usually expressed in electron volts. Thus, the electron affinity of the hydrogen atom is 0.75 eV, oxygen - 1.47 eV, fluorine - 3.52 eV.

The electron affinities of metal atoms are typically close to zero or negative; It follows from this that for atoms of most metals the addition of electrons is energetically unfavorable. The electron affinity of nonmetal atoms is always positive and the greater, the closer the nonmetal is located to the noble gas in the periodic table; this indicates an increase in non-metallic properties as the end of the period approaches.

There are many repeating sequences in nature:

Seasons;
Times of Day;
days of the week…

In the mid-19th century, D.I. Mendeleev noticed that the chemical properties of elements also have a certain sequence (they say that this idea came to him in a dream). The result of the scientist’s wonderful dreams was the Periodic Table of Chemical Elements, in which D.I. Mendeleev arranged chemical elements in order of increasing atomic mass. In the modern table, chemical elements are arranged in ascending order of the element's atomic number (the number of protons in the nucleus of an atom).

The atomic number is shown above the symbol of a chemical element, below the symbol is its atomic mass (the sum of protons and neutrons). Please note that the atomic mass of some elements is not a whole number! Remember isotopes! Atomic mass is the weighted average of all isotopes of an element found in nature under natural conditions.

Below the table are lanthanides and actinides.

Metals, non-metals, metalloids

Located in the Periodic Table to the left of a stepped diagonal line that begins with Boron (B) and ends with polonium (Po) (the exceptions are germanium (Ge) and antimony (Sb). It is easy to see that metals occupy most of the Periodic Table. Basic properties of metals : hard (except mercury); shiny; good electrical and thermal conductors; plastic; malleable; easily give up electrons.

The elements located to the right of the B-Po stepped diagonal are called non-metals. The properties of non-metals are exactly the opposite of those of metals: poor conductors of heat and electricity; fragile; non-malleable; non-plastic; usually accept electrons.

Metalloids

Between metals and non-metals there are semimetals(metalloids). They are characterized by the properties of both metals and non-metals. Semimetals have found their main application in industry in the production of semiconductors, without which not a single modern microcircuit or microprocessor is conceivable.

Periods and groups

As mentioned above, the periodic table consists of seven periods. In each period, the atomic numbers of elements increase from left to right.

The properties of elements change sequentially in periods: thus sodium (Na) and magnesium (Mg), located at the beginning of the third period, give up electrons (Na gives up one electron: 1s 2 2s 2 2p 6 3s 1 ; Mg gives up two electrons: 1s 2 2s 2 2p 6 3s 2). But chlorine (Cl), located at the end of the period, takes one element: 1s 2 2s 2 2p 6 3s 2 3p 5.

In groups, on the contrary, all elements have the same properties. For example, in group IA(1), all elements from lithium (Li) to francium (Fr) donate one electron. And all elements of group VIIA(17) take one element.

Some groups are so important that they have received special names. These groups are discussed below.

Group IA(1). Atoms of elements of this group have only one electron in their outer electron layer, so they easily give up one electron.

The most important alkali metals are sodium (Na) and potassium (K), since they play an important role in human life and are part of salts.

Electronic configurations:

Li- 1s 2 2s 1 ;
Na- 1s 2 2s 2 2p 6 3s 1 ;
K- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1

Group IIA(2). Atoms of elements of this group have two electrons in their outer electron layer, which they also give up during chemical reactions. The most important element is calcium (Ca) - the basis of bones and teeth.

Electronic configurations:

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

Group VIIA(17). Atoms of elements of this group usually receive one electron each, because There are five elements on the outer electronic layer and one electron is just missing from the “complete set”.

The most well-known elements of this group: chlorine (Cl) - is part of salt and bleach; Iodine (I) is an element that plays an important role in the activity of the human thyroid gland.

Electronic Configuration:

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 4p 5

Group VIII(18). Atoms of elements of this group have a fully “complete” outer electron layer. Therefore, they “don’t” need to accept electrons. And they “don’t want” to give them away. Hence, the elements of this group are very “reluctant” to enter into chemical reactions. For a long time it was believed that they do not react at all (hence the name “inert”, i.e. “inactive”). But chemist Neil Bartlett discovered that some of these gases can still react with other elements under certain conditions.

Electronic configurations:

Ne- 1s 2 2s 2 2p 6 ;
Ar- 1s 2 2s 2 2p 6 3s 2 3p 6 ;
Kr- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6

Valence elements in groups

It is easy to notice that within each group the elements are similar to each other in their valence electrons (electrons of s and p orbitals located on the outer energy level).

Alkali metals have 1 valence electron:

Li- 1s 2 2s 1 ;
Na- 1s 2 2s 2 2p 6 3s 1 ;
K- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1

Alkaline earth metals have 2 valence electrons:

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

Halogens have 7 valence electrons:

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 4p 5

Inert gases have 8 valence electrons:

Ne- 1s 2 2s 2 2p 6 ;
Ar- 1s 2 2s 2 2p 6 3s 2 3p 6 ;
Kr- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6

For more information, see the article Valency and the Table of Electronic Configurations of Atoms of Chemical Elements by Period.

Let us now turn our attention to the elements located in groups with symbols IN. They are located in the center of the periodic table and are called transition metals.

A distinctive feature of these elements is the presence in the atoms of electrons that fill d-orbitals:

Sc- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 ;
Ti- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2

Separately from the main table are located lanthanides And actinides- these are the so-called internal transition metals. In the atoms of these elements, electrons fill f-orbitals:

Ce- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 4d 10 5s 2 5p 6 4f 1 5d 1 6s 2 ;
Th- 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 4d 10 5s 2 5p 6 4f 14 5d 10 6s 2 6p 6 6d 2 7s 2

PERIODIC SYSTEM, an ordered set of chemicals. elements, their natures. , which is a table expression. The prototype of the periodic chemical systems elements was based on the table “Experience of a system of elements based on their chemical similarity,” compiled by D. I. Mendeleev on March 1, 1869 (Fig. 1). Lastly Over the years, the scientist improved the table, developed ideas about periods and groups of elements and the place of an element in the system. In 1870, Mendeleev called the system natural, and in 1871 periodic. As a result, even then the periodic system in many respects acquired its modern form. structural outlines. Based on it, Mendeleev predicted the existence of saints ca. 10 unknown elements; these predictions were subsequently confirmed.

Rice. 1 Table “Experience of a system of elements based on their chemical similarity” (D. I. Mendeleev. I myrtle 1869).

However, over the next more than 40 years, the periodic table means. degree was only empirical. generalization of facts, since there was no physical explanation of reasons periodic. changes in CB-B elements depending on their increase. Such an explanation was impossible without well-founded ideas about the structure (see). Therefore, the most important milestone in the development of the periodic table was the planetary (nuclear) model proposed by E. Rutherford (1911). In 1913, A. van den Broek came to the conclusion that an element in the periodic table is numerically equal to posi. charge (Z) of its nucleus. This conclusion was experimentally confirmed by G. Moseley (Moseley's law, 1913-14). As a result, periodic the law received strict physical formulation, it was possible to unambiguously determine the following. boundary of the periodic table (H as an element with a minimum of Z=1), estimate the exact number of elements between H and U and determine which elements have not yet been discovered (Z = 43, 61, 72, 75, 85, 87). The theory of the periodic table was developed in the beginning. 1920s (see below).

Structure of the periodic table. The modern periodic system includes 109 chemical elements (there is information about the synthesis in 1988 of an element with Z = 110). Of these in natural objects found 89; all elements following U, or (Z = 93 109), as well as Tc (Z = 43), Pm (Z = 61) and At (Z = 85) were artificially synthesized using decomp. . Elements with Z = 106 109 have not yet received names, so there are no corresponding symbols in the tables; for an element with Z = 109 the maximum values are still unknown. long-lived

Over the entire history of the periodic table, more than 500 different versions of its image have been published. This was due to attempts to find a rational solution to certain controversial problems of the structure of the periodic table (placement of H, lanthanides, etc.). Naib. spread as follows. tabular forms of expression of the periodic system: 1) the short one was proposed by Mendeleev (in its current form it is placed at the beginning of the volume on the colored flyleaf); 2) the long one was developed by Mendeleev, improved in 1905 by A. Werner (Fig. 2); 3) staircase published in 1921 H. (Fig. 3). In recent decades, short and long forms have been especially widely used, as they are visual and practically convenient. All listed. forms have certain advantages and disadvantages. However, it is hardly possible to offer k.-l. univers. variant of the representation of the periodic table, which would adequately reflect all the diversity of the world of chemistry. elements and the specifics of changes in their chemical. behavior as Z increases.

Fundam. The principle of constructing the periodic table is to distinguish periods (horizontal rows) and groups (vertical columns) of elements in it. The modern periodic system consists of 7 periods (the seventh, not yet completed, should end with a hypothetical element with Z = 118) and 8 groups. The period is called. a set of elements that begins (or the first period) and ends. The numbers of elements in periods naturally increase and, starting from the second, repeat in pairs: 8, 8, 18, 18, 32, 32, ... (a special case is the first period, containing only two elements). The group of elements does not have a clear definition; Formally, its number corresponds to max. the meaning of its constituent elements, but this condition is not met in a number of cases. Each group is divided into main (a) and secondary (b) subgroups; each of them contains elements that are chemically similar. St. you, which are characterized by the same external structure. electronic shells. In most groups, elements of subgroups a and b exhibit a certain chemical. similarity, prem. in higher .

Group VIII occupies a special place in the structure of the periodic table. For a long time time, only elements of the “triads” were attributed to it: Fe-Co-Ni and (Ru Rh Pd and Os-Ir-Pt), and all were placed in independent positions. zero group; therefore, the periodic table contained 9 groups. After in the 60s. were received conn. Xe, Kr and Rn began to be placed in subgroup VIIIa, and the zero group was abolished. The elements of the triads made up subgroup VIII6. This “structural design” of group VIII now appears in almost all published expressions of the periodic table.

Will distinguish. The feature of the first period is that it contains only 2 elements: H and He. due to the holy - unities. an element that does not have a clearly defined place in the periodic table. The symbol H is placed either in subgroup Ia, or in subgroup VIIa, or in both at the same time, enclosing the symbol in brackets in one of the subgroups, or, finally, depicting it as separated. fonts. These ways of arranging H are based on the fact that it has certain formal similarities with both .

Rice. 2. Long form periodic. chemical systems elements (modern version). Rice. 3. Ladder form periodic. chemical systems elements (H., 1921).

The second period (Li-Ne), containing 8 elements, begins with Li (unities, + 1); followed by Be(+2). Metallic character B (+3) is weakly expressed, and the next one, C, is typical (+4). The following are N, O, F and Ne-non-metals, with only N having the highest + 5 corresponding to the group number; O and F are among the most active.

The third period (Na-Ar) also includes 8 elements, the nature of the chemical change. St. in which is in many ways similar to that observed in the second period. However, Mg and Al are more “metallic” than the corresponding ones. Be and B. The remaining elements are Si, P, S, Cl and Ar non-metals; they all exhibit , equal to the group number, except Ar. T.arr., in the second and third periods, as Z increases, a weakening of the metallic and an increase in non-metallic is observed. nature of the elements.

All elements of the first three periods belong to subgroups a. According to modern terminology, elements belonging to subgroups Ia and IIa are called. I-elements (in the color table their symbols are given in red), to subgroups IIIa-VIIIa-p-elements (orange symbols).

The fourth period (K-Kr) contains 18 elements. After K and alkaline-earth. Ca (s-elements) follows a series of 10 so-called. transition (Sc-Zn), or d-elements (blue symbols), which are included in subgroups b. The majority (all of them - ) exhibit higher , equal to the group number, excluding the Fe-Co-Ni triad, where Fe under certain conditions has +6, and Co and Ni are maximally trivalent. Elements from Ga to Kr belong to subgroups a (p-elements), and the nature of the change in their properties is in many ways similar to the change in the properties of elements of the second and third periods in the corresponding intervals of Z values. For Kr, several were obtained. relatively stable compounds, mainly with F.

The fifth period (Rb-Xe) is constructed similarly to the fourth; it also has an insert of 10 transition, or d-elements (Y-Cd). Peculiarities of changes in the strength of elements in the period: 1) in the triad Ru-Rh-Pd shows a maximum of 4-8; 2) all elements of subgroups a, including Xe, exhibit higher values equal to the group number; 3) I has weak metallic properties. St. T. example, the properties of the elements of the fourth and fifth periods change more complexly as Z increases than the properties of the elements in the second and third periods, which is primarily due to the presence of transition d-elements.

The sixth period (Cs-Rn) contains 32 elements. In addition to ten d-elements (La, Hf-Hg), it includes a family of 14 f-elements (black symbols, from Ce to Lu)-lanthanides. They are very similar in chemistry. Holy to you (preferably at +3) and therefore cannot. placed according to different system groups. In the short form of the periodic table, all lanthanides are included in subgroup IIIa (La), and their totality is deciphered below the table. This technique is not without its drawbacks, since the 14 elements appear to be outside the system. In the long and ladder forms of the periodic system, the specificity is reflected in the general background of its structure. Dr. features of the elements of the period: 1) in the Os Ir Pt triad, only Os exhibits a max. +8; 2) At has a more pronounced metallic effect compared to I. character; 3) Rn max. is reactive from, but its strong chemistry makes it difficult to study. St.

The seventh period, like the sixth, should contain 32 elements, but is not yet completed. Fr and Ra elements respectively. subgroups Ia and IIa, Ac is an analogue of elements of subgroup III6. According to the actinide concept of G. Seaborg (1944), after Ac comes a family of 14 f elements (Z = 90 103). In the short form of the periodic table, the latter are included in Ac and are similarly written as dept. line below the table. This technique assumed the presence of a certain chemical. similarities between elements of two f-families. However, a detailed study showed that they exhibit a much wider range, including such as +7 (Np, Pu, Am). In addition, heavy ones are characterized by stabilization of lower ones (+ 2 or even +1 for Md).

Chemical assessment nature of Ku (Z = 104) and Ns (Z = 105), synthesized in a number of single, very short-lived ones, allowed us to conclude that these elements are analogues of respectively. Hf and Ta, i.e. d-elements, and should be located in subgroups IV6 and V6. Chem. elements with Z = 106 109 were not carried out, but it can be assumed that they belong to the seventh period. Computer calculations indicate that elements with Z = 113,118 belong to p-elements (subgroup IIIa VIIIa).

Theory of the periodic table was preem. created by H. (1913 21) on the basis of the quantum model he proposed. Taking into account the specifics of changes in the properties of elements in the periodic system and information about them, he developed a scheme for constructing electronic configurations as Z increases, making it the basis for explaining the phenomenon of periodicity and the structure of the periodic system. This scheme is based on a certain sequence of filling shells (also called layers, levels) and subshells (shells, sublevels) in accordance with the increase in Z. Similar electronic configurations ext. electron shells repeat periodically, which determines the periodicity. chemical change St. elements. This is what ch. cause physical the nature of the phenomenon of periodicity. Electronic shells, with the exception of those that correspond to values 1 and 2 of the main quantum number l, are not filled sequentially and monotonically until their complete completion (the numbers in the sequential shells are: 2, 8, 18, 32, 50,... ); their construction is periodically interrupted by the appearance of aggregates (constituting certain subshells), which correspond to large values of n. This is the essence of beings. peculiarity of the “electronic” interpretation of the structure of the periodic table.

The scheme for the formation of electronic configurations, which underlies the theory of the periodic system, thus reflects a certain sequence of appearance as Z grows of aggregates (subshells), characterized by certain values of the principal and orbital (l) quantum numbers. This scheme is generally written in the form of a table. (see below).

Vertical lines separate the subshells, which are filled into the elements that make up the sequence. periods of the periodic system (period numbers are indicated by numbers at the top); Subshells that complete the formation of shells with a given item are highlighted in bold.

The numbers in shells and subshells are defined on . In relation to , as particles with a half-integer, he postulates that in no way. two with identical values of all quantum numbers. The capacities of the shells and subshells are equal, respectively. 2n 2 and 2(2l + 1). This principle does not define.


Period	1	2	3	4	5	6	7
Electronic configuration	1s	2s 2р	3s 3р	4s 3d 4p	5s 4d 5p	6s 4f 5d 6p	7s 5f 6d 7p
n	l	22	33	434	545	6456	7567
l	0	01	01	021	021	0321	0321
	2	26	26	2106	2106	214106	214106
Number of elements in period	2	8	8	18	18	32	32

however, the sequence of formation of electronic configurations as Z increases. From the above diagram, the capacitances are found in series. periods: 2, 8, 18, 32, 32, ....

Each period begins with an element in which it first appears with a given value of n at l = 0 (ns 1 -elements), and ends with an element in which a subshell with the same n and l = 1 is filled (np 6 -elements You); exception is the first period (1s elements only). All s- and p-elements belong to subgroups a. Subgroups b include elements in which shells that were previously left unfinished are completed (the values of h are less than the period number, l = 2 and 3). The first three periods include elements only of subgroups a, i.e. s- and p-elements.

The actual scheme for constructing electronic configurations is described by the so-called. (n + l)-rule formulated (1951) by V. M. Klechkovsky. The construction of electronic configurations occurs in accordance with the successive increase in the sum (n + /). Moreover, within each such sum, subshells with larger l and smaller n are first filled, then with smaller l and larger n.

Starting from the sixth period, the construction of electronic configurations actually becomes more complex, which is expressed in the violation of clear boundaries between successively filled subshells. For example, the 4f electron appears not in La with Z = 57, but in the next one Ce (Z = 58); sequential the construction of the 4f subshell is interrupted in Gd (Z = 64, presence of a 5d electron). Such a “blurring of periodicity” clearly affects the seventh period for Z > 89, which is reflected in the properties of the elements.

The real scheme was not originally derived from the k.-l. rigorous theoretical representations. It was based on well-known chem. holy elements and information about their spectra. Valid physical the real scheme received its justification through the application of methods to the description of the structure. In quantum mech. interpretation of the theory of structure, the concept of electronic shells and subshells with a strict approach has lost its original meaning; the concept of atomic is now widely used. Nevertheless, the developed principle of physical interpretation of the phenomenon of periodicity has not lost its significance and, to a first approximation, explains the theoretical theory quite comprehensively. basics of the periodic table. In any case, the published forms of the periodic table reflect the idea of the nature of the distribution among shells and subshells.

Structure and chemical properties of elements. Main features of chemistry. the behavior of elements is determined by the nature of the configurations of the outer (one or two) electron shells. These features are different for elements of subgroups a (s- and p-elements), subgroups b (d-elements), f-families ( and ).

A special place is occupied by the 1s elements of the first period (H and He). due to the presence in only one there is a large differenceSt. The configuration of He (1s 2) is exceptional, which determines its chemical inertia. Since the elements of subgroups a are filled with ext. electron shells (with n equal to the period number), the properties of the elements change noticeably as Z increases in the corresponding periods, which is expressed in the weakening of metallic and the strengthening of non-metallic. St. All except H and He are p-elements. At the same time, in each subgroup a, as Z increases, an increase in metallicity is observed. St. These patterns are explained by the weakening of the external binding energy. with the core during the transition from period to period.

Meaning of the periodic table. This system has played and continues to play a huge role in the development of pluralism. natural science disciplines. She became an important link in the atomic pier. teachings, contributed to the formulation of modern. the concept of "chemical element" and the clarification of ideas about simple substances and compounds. influence on the development of the theory of structure and the emergence of the concept of isotopy. Strictly scientific is connected with the periodic system. formulation of the forecasting problem in thatmanifested itself both in the prediction of the existence of unknown elements and their properties, as well as new chemical features. behavior of already opened elements. The periodic table is the most important basis of inorg. ; it serves, for example, the tasks of synthesizing materials with predetermined properties, creating new materials, in particular semiconductor materials, and selecting specific materials. for diff. chem. processes. Periodic system - scientific. general and non-organizational teaching base , as well as certain branches of atomic physics.

Lit.: Mendeleev D.I., Periodic law. Basic articles, M., 1958; Kedrov B. M.. Three aspects of atomism, part 3. Mendeleev’s Law, M., 1969; Trifonov D N., On the quantitative interpretation of periodicity, M., 1971; Trifonov D. N., Krivomazov A. N., Lisnevsky Yu. I., The doctrine of periodicity and the doctrine of. Combined chronology of the most important events. M., 1974; Karapetyami MX. Drakii S.I., Stroenie, M., 1978; The doctrine of periodicity. History and modernity. Sat. articles. M.. 1981. Korolkov D.V., Fundamentals, M., 1982; Melnikov V.P., Dmitriev I.S. Additional types of periodicity in the periodic system of D.I. Mendeleev, M. 1988. D.N Trifonov.

The properties of chemical elements make it possible to combine them into appropriate groups. On this principle, the periodic system was created, which changed the idea of existing substances and made it possible to assume the existence of new, previously unknown elements.

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Mendeleev's periodic table

The periodic table of chemical elements was compiled by D.I. Mendeleev in the second half of the 19th century. What is it and what is it for? It unites all chemical elements in order of increasing atomic weight, and they are all arranged in such a way that their properties change in a periodic manner.

Mendeleev's periodic system brought together into a single system all existing elements, previously considered simply separate substances.

Based on its study, new chemical substances were predicted and subsequently synthesized. The significance of this discovery for science cannot be overestimated, it was significantly ahead of its time and gave impetus to the development of chemistry for many decades.

There are three most common table options, which are conventionally called “short”, “long” and “extra-long” ». The main table is considered to be a long table, it officially approved. The difference between them is the arrangement of elements and the length of periods.

What is a period

The system contains 7 periods. They are presented graphically as horizontal lines. In this case, a period can have one or two lines, called rows. Each subsequent element differs from the previous one by increasing the nuclear charge (number of electrons) by one.

To keep it simple, a period is a horizontal row of the periodic table. Each of them begins with metal and ends with an inert gas. Actually, this creates periodicity - the properties of elements change within one period, repeating again in the next. The first, second and third periods are incomplete, they are called small and contain 2, 8 and 8 elements, respectively. The rest are complete, they have 18 elements each.

What is a group

A group is a vertical column, containing elements with the same electronic structure or, more simply, with the same higher value. The officially approved long table contains 18 groups, which begin with alkali metals and end with noble gases.

Each group has its own name, making it easier to search or classify elements. Metallic properties are enhanced, regardless of the element, from top to bottom. This is due to an increase in the number of atomic orbits - the more there are, the weaker the electronic bonds, which makes the crystal lattice more pronounced.

Metals in the periodic table

Metals in the table Mendeleev have a predominant number, their list is quite extensive. They are characterized by common characteristics; they are heterogeneous in their properties and are divided into groups. Some of them have little in common with metals in the physical sense, while others can exist only for a fraction of a second and are absolutely not found in nature (at least on the planet), since they were created, or rather, calculated and confirmed in laboratory conditions, artificially. Each group has its own characteristics, the name is quite noticeably different from the others. This difference is especially pronounced in the first group.

Position of metals

What is the position of metals in the periodic table? Elements are arranged by increasing atomic mass, or number of electrons and protons. Their properties change periodically, so there is no neat placement on a one-to-one basis in the table. How to identify metals, and is it possible to do this using the periodic table? In order to simplify the question, a special technique was invented: conditionally, a diagonal line is drawn from Bor to Polonius (or to Astatus) at the junctions of the elements. Those on the left are metals, those on the right are non-metals. This would be very simple and cool, but there are exceptions - Germanium and Antimony.

This “methodology” is a kind of cheat sheet; it was invented only to simplify the memorization process. For a more accurate representation, it should be remembered that the list of nonmetals is only 22 elements, therefore, answering the question, how many metals are contained in the periodic table?

In the figure you can clearly see which elements are non-metals and how they are arranged in the table by groups and periods.

General physical properties

There are general physical properties of metals. These include:

Plastic.
Characteristic shine.
Electrical conductivity.
High thermal conductivity.
All except mercury are in a solid state.

It should be understood that the properties of metals vary greatly regarding their chemical or physical essence. Some of them bear little resemblance to metals in the ordinary sense of the term. For example, mercury occupies a special position. Under normal conditions, it is in a liquid state and does not have a crystal lattice, the presence of which other metals owe their properties to. The properties of the latter in this case are conditional; mercury is similar to them to a greater extent in its chemical characteristics.

Interesting! Elements of the first group, alkali metals, are not found in pure form, but are found in various compounds.

The softest metal existing in nature, cesium, belongs to this group. It, like other alkaline substances, has little in common with more typical metals. Some sources claim that in fact, the softest metal is potassium, which is difficult to dispute or confirm, since neither one nor the other element exists on its own - when released as a result of a chemical reaction, they quickly oxidize or react.

The second group of metals - alkaline earth metals - are much closer to the main groups. The name "alkaline earth" comes from ancient times, when oxides were called "earths" because they had a loose, crumbly structure. Metals starting from group 3 have more or less familiar (in the everyday sense) properties. As the group number increases, the amount of metals decreases, being replaced by non-metallic elements. The last group consists of inert (or noble) gases.

Determination of metals and non-metals in the periodic table. Simple and complex substances.

Simple substances (metals and non-metals)

Conclusion

The ratio of metals and non-metals in the periodic table clearly weighs in favor of the former. This situation indicates that the group of metals is combined too broadly and requires a more detailed classification, which is recognized by the scientific community.

The periodic system is an ordered set of chemical elements, their natural classification, which is a graphic (tabular) expression of the periodic law of chemical elements. Its structure, in many ways similar to the modern one, was developed by D. I. Mendeleev on the basis of the periodic law in 1869–1871.

The prototype of the periodic system was the “Experience of a system of elements based on their atomic weight and chemical similarity”, compiled by D. I. Mendeleev on March 1, 1869. Over the course of two and a half years, the scientist continuously improved the “Experience of a System”, introduced the idea of groups, series and periods of elements. As a result, the structure of the periodic table acquired largely modern outlines.

The concept of the place of an element in the system, determined by the numbers of the group and period, became important for its evolution. Based on this concept, Mendeleev came to the conclusion that it was necessary to change the atomic masses of some elements: uranium, indium, cerium and its satellites. This was the first practical application of the periodic table. Mendeleev also predicted for the first time the existence and properties of several unknown elements. The scientist described in detail the most important properties of eka-aluminium (the future of gallium), eka-boron (scandium) and eka-silicon (germanium). In addition, he predicted the existence of analogues of manganese (future technetium and rhenium), tellurium (polonium), iodine (astatine), cesium (France), barium (radium), tantalum (protactinium). The scientist's predictions regarding these elements were of a general nature, since these elements were located in little-studied areas of the periodic table.

The first versions of the periodic system largely represented only an empirical generalization. After all, the physical meaning of the periodic law was unclear; there was no explanation for the reasons for the periodic change in the properties of elements depending on the increase in atomic masses. In this regard, many problems remained unresolved. Are there boundaries of the periodic table? Is it possible to determine the exact number of existing elements? The structure of the sixth period remained unclear - what was the exact amount of rare earth elements? It was unknown whether elements between hydrogen and lithium still existed, what the structure of the first period was. Therefore, right up to the physical substantiation of the periodic law and the development of the theory of the periodic system, serious difficulties arose more than once. The discovery in 1894–1898 was unexpected. five inert gases that seemed to have no place in the periodic table. This difficulty was eliminated thanks to the idea of including an independent zero group in the structure of the periodic table. Mass discovery of radioelements at the turn of the 19th and 20th centuries. (by 1910 their number was about 40) led to a sharp contradiction between the need to place them in the periodic table and its existing structure. There were only 7 vacancies for them in the sixth and seventh periods. This problem was solved by the establishment of shift rules and the discovery of isotopes.

One of the main reasons for the impossibility of explaining the physical meaning of the periodic law and the structure of the periodic system was that it was unknown how the atom was structured (see Atom). The most important milestone in the development of the periodic table was the creation of the atomic model by E. Rutherford (1911). On its basis, the Dutch scientist A. Van den Broek (1913) suggested that the serial number of an element in the periodic table is numerically equal to the charge of the nucleus of its atom (Z). This was experimentally confirmed by the English scientist G. Moseley (1913). The periodic law received a physical justification: the periodicity of changes in the properties of elements began to be considered depending on the Z - charge of the nucleus of the element's atom, and not on the atomic mass (see Periodic law of chemical elements).

As a result, the structure of the periodic table was significantly strengthened. The lower limit of the system has been determined. This is hydrogen - the element with a minimum Z = 1. It has become possible to accurately estimate the number of elements between hydrogen and uranium. “Gaps” in the periodic table were identified, corresponding to unknown elements with Z = 43, 61, 72, 75, 85, 87. However, questions about the exact number of rare earth elements remained unclear and, most importantly, the reasons for the periodicity of changes in the properties of elements were not revealed depending on Z.

Based on the established structure of the periodic system and the results of studying atomic spectra, the Danish scientist N. Bohr in 1918–1921. developed ideas about the sequence of construction of electronic shells and subshells in atoms. The scientist came to the conclusion that similar types of electronic configurations of the outer shells of atoms are periodically repeated. Thus, it was shown that the periodicity of changes in the properties of chemical elements is explained by the existence of periodicity in the construction of electronic shells and subshells of atoms.

The periodic table covers more than 100 elements. Of these, all transuranium elements (Z = 93–110), as well as elements with Z = 43 (technetium), 61 (promethium), 85 (astatine), 87 (france) were obtained artificially. Over the entire history of the existence of the periodic system, a very large number (>500) of variants of its graphic representation have been proposed, mainly in the form of tables, but also in the form of various geometric figures (spatial and planar), analytical curves (spirals, etc.), etc. The most widespread are short, semi-long, long and ladder forms of tables. Currently, short form is preferred.

The fundamental principle of constructing the periodic table is its division into groups and periods. Mendeleev's concept of series of elements is not used today, since it is devoid of physical meaning. The groups, in turn, are divided into main (a) and secondary (b) subgroups. Each subgroup contains elements - chemical analogues. Elements of the a- and b-subgroups in most groups also show a certain similarity with each other, mainly in higher oxidation states, which, as a rule, are equal to the group number. A period is a collection of elements that begins with an alkali metal and ends with an inert gas (a special case is the first period). Each period contains a strictly defined number of elements. The periodic table consists of eight groups and seven periods, with the seventh period not yet completed.

Peculiarity first period is that it contains only 2 gaseous elements in free form: hydrogen and helium. The place of hydrogen in the system is ambiguous. Since it exhibits properties common to alkali metals and halogens, it is placed either in the 1a-, or in the Vlla-subgroup, or in both at the same time, enclosing the symbol in brackets in one of the subgroups. Helium is the first representative of the VIIIa‑subgroup. For a long time, helium and all inert gases were separated into an independent zero group. This position required revision after the synthesis of the chemical compounds krypton, xenon and radon. As a result, the noble gases and elements of the former Group VIII (iron, cobalt, nickel and platinum metals) were combined within one group.

Second the period contains 8 elements. It begins with the alkali metal lithium, whose only oxidation state is +1. Next comes beryllium (metal, oxidation state +2). Boron already exhibits a weakly expressed metallic character and is a non-metal (oxidation state +3). Next to boron, carbon is a typical nonmetal that exhibits both +4 and −4 oxidation states. Nitrogen, oxygen, fluorine and neon are all non-metals, with nitrogen having the highest oxidation state of +5 corresponding to the group number. Oxygen and fluorine are among the most active nonmetals. The inert gas neon ends the period.

Third period (sodium - argon) also contains 8 elements. The nature of the change in their properties is largely similar to that observed for elements of the second period. But there is also some specificity here. Thus, magnesium, unlike beryllium, is more metallic, as is aluminum compared to boron. Silicon, phosphorus, sulfur, chlorine, argon are all typical non-metals. And all of them, except argon, exhibit higher oxidation states equal to the group number.

As we can see, in both periods, as Z increases, there is a clear weakening of the metallic and strengthening of the nonmetallic properties of the elements. D.I. Mendeleev called the elements of the second and third periods (in his words, small) typical. Elements of small periods are among the most common in nature. Carbon, nitrogen and oxygen (along with hydrogen) are organogens, i.e. the main elements of organic matter.

All elements of the first - third periods are placed in a-subgroups.

Fourth period (potassium - krypton) contains 18 elements. According to Mendeleev, this is the first big period. After the alkali metal potassium and the alkaline earth metal calcium comes a series of elements consisting of 10 so-called transition metals (scandium - zinc). All of them are included in b-subgroups. Most transition metals exhibit higher oxidation states equal to the group number, except iron, cobalt and nickel. The elements, from gallium to krypton, belong to the a-subgroups. A number of chemical compounds are known for krypton.

Fifth The period (rubidium - xenon) is similar in structure to the fourth. It also contains an insert of 10 transition metals (yttrium - cadmium). The elements of this period have their own characteristics. In the triad ruthenium - rhodium - palladium, compounds are known for ruthenium where it exhibits an oxidation state of +8. All elements of a-subgroups exhibit higher oxidation states equal to the group number. The features of changes in properties of elements of the fourth and fifth periods as Z increases are more complex in comparison with the second and third periods.

Sixth period (cesium - radon) includes 32 elements. This period, in addition to 10 transition metals (lanthanum, hafnium - mercury), also contains a set of 14 lanthanides - from cerium to lutetium. Elements from cerium to lutetium are chemically very similar, and for this reason they have long been included in the family of rare earth elements. In the short form of the periodic table, a series of lanthanides is included in the lanthanum cell, and the decoding of this series is given at the bottom of the table (see Lanthanides).

What is the specificity of the elements of the sixth period? In the triad osmium - iridium - platinum, the oxidation state of +8 is known for osmium. Astatine has a fairly pronounced metallic character. Radon has the greatest reactivity of all noble gases. Unfortunately, due to the fact that it is highly radioactive, its chemistry has been little studied (see Radioactive elements).

Seventh the period starts from France. Like the sixth, it should also contain 32 elements, but 24 of them are still known. Francium and radium are respectively elements of the Ia and IIa subgroups, actinium belongs to the IIIb subgroup. Next comes the actinide family, which includes elements from thorium to lawrencium and is placed similarly to the lanthanides. The decoding of this series of elements is also given at the bottom of the table.

Now let's see how the properties of chemical elements change in subgroups periodic system. The main pattern of this change is the strengthening of the metallic character of the elements as Z increases. This pattern is especially clearly manifested in the IIIa–VIIa subgroups. For metals of Ia–IIIa subgroups, an increase in chemical activity is observed. For elements of IVa–VIIa subgroups, as Z increases, a weakening of the chemical activity of the elements is observed. For b-subgroup elements, the nature of the change in chemical activity is more complex.

The theory of the periodic system was developed by N. Bohr and other scientists in the 20s. XX century and is based on a real scheme for the formation of electronic configurations of atoms (see Atom). According to this theory, as Z increases, the filling of electron shells and subshells in the atoms of elements included in the periods of the periodic table occurs in the following sequence:

Period numbers
1	2	3	4	5	6	7
1s	2s2p	3s3p	4s3d4p	5s4d5p	6s4f5d6p	7s5f6d7p

Based on the theory of the periodic system, we can give the following definition of a period: a period is a set of elements starting with an element with a value n equal to the period number and l = 0 (s-elements) and ending with an element with the same value n and l = 1 (p-elements elements) (see Atom). The exception is the first period, which contains only 1s elements. From the theory of the periodic system, the numbers of elements in periods follow: 2, 8, 8, 18, 18, 32...

In the table, the symbols of elements of each type (s-, p-, d- and f-elements) are depicted on a specific color background: s-elements - on red, p-elements - on orange, d-elements - on blue, f-elements - on green. Each cell shows the atomic numbers and atomic masses of the elements, as well as the electronic configurations of the outer electron shells.

From the theory of the periodic system it follows that the a-subgroups include elements with n equal to the period number, and l = 0 and 1. The b-subgroups include those elements in the atoms of which the completion of shells that previously remained incomplete occurs. That is why the first, second and third periods do not contain elements of b-subgroups.

The structure of the periodic table of elements is closely related to the structure of atoms of chemical elements. As Z increases, similar types of configuration of the outer electron shells periodically repeat. Namely, they determine the main features of the chemical behavior of elements. These features manifest themselves differently for elements of the a-subgroups (s- and p-elements), for elements of the b-subgroups (transition d-elements) and elements of the f-families - lanthanides and actinides. A special case is represented by the elements of the first period - hydrogen and helium. Hydrogen is characterized by high chemical activity because its only 1s electron is easily removed. At the same time, the configuration of helium (1s 2) is very stable, which determines its chemical inactivity.

For elements of the a-subgroups, the outer electron shells of the atoms are filled (with n equal to the period number), so the properties of these elements change noticeably as Z increases. Thus, in the second period, lithium (2s configuration) is an active metal that easily loses its only valence electron ; beryllium (2s 2) is also a metal, but less active due to the fact that its outer electrons are more tightly bound to the nucleus. Further, boron (2s 2 p) has a weakly expressed metallic character, and all subsequent elements of the second period, in which the 2p subshell is built, are already non-metals. The eight-electron configuration of the outer electron shell of neon (2s 2 p 6) - an inert gas - is very strong.

The chemical properties of elements of the second period are explained by the desire of their atoms to acquire the electronic configuration of the nearest inert gas (helium configuration for elements from lithium to carbon or neon configuration for elements from carbon to fluorine). This is why, for example, oxygen cannot exhibit a higher oxidation state equal to its group number: it is easier for it to achieve the neon configuration by acquiring additional electrons. The same nature of changes in properties manifests itself in the elements of the third period and in the s- and p-elements of all subsequent periods. At the same time, the weakening of the strength of the bond between outer electrons and the nucleus in a-subgroups as Z increases is manifested in the properties of the corresponding elements. Thus, for s‑elements there is a noticeable increase in chemical activity as Z increases, and for p‑elements there is an increase in metallic properties.

In the atoms of transition d‑elements, previously incomplete shells are completed with the value of the main quantum number n, one less than the period number. With a few exceptions, the configuration of the outer electron shells of the atoms of transition elements is ns 2. Therefore, all d-elements are metals, and that is why the changes in the properties of d-elements as Z increases are not as dramatic as those observed for s- and p-elements. In higher oxidation states, d-elements show a certain similarity with p-elements of the corresponding groups of the periodic table.

The peculiarities of the properties of the elements of triads (VIIIb-subgroup) are explained by the fact that the b-subshells are close to completion. This is why iron, cobalt, nickel and platinum metals, as a rule, do not tend to produce compounds in higher oxidation states. The only exceptions are ruthenium and osmium, which give the oxides RuO 4 and OsO 4 . For elements of subgroups Ib and IIb, the d-subshell is actually complete. Therefore, they exhibit oxidation states equal to the group number.

In the atoms of lanthanides and actinides (all of them are metals), previously incomplete electron shells are completed with the value of the main quantum number n being two units less than the period number. In the atoms of these elements, the configuration of the outer electron shell (ns 2) remains unchanged, and the third outer N‑shell is filled with 4f‑electrons. This is why the lanthanides are so similar.

For actinides the situation is more complicated. In atoms of elements with Z = 90–95, the 6d and 5f electrons can take part in chemical interactions. Therefore, actinides have many more oxidation states. For example, for neptunium, plutonium and americium, compounds are known where these elements appear in the heptavalent state. Only for elements, starting with curium (Z = 96), the trivalent state becomes stable, but this also has its own characteristics. Thus, the properties of the actinides differ significantly from the properties of the lanthanides, and the two families therefore cannot be considered similar.

The actinide family ends with the element with Z = 103 (lawrencium). An assessment of the chemical properties of kurchatovium (Z = 104) and nilsborium (Z = 105) shows that these elements should be analogues of hafnium and tantalum, respectively. Therefore, scientists believe that after the actinide family in atoms, the systematic filling of the 6d subshell begins. The chemical nature of elements with Z = 106–110 has not been assessed experimentally.

The final number of elements that the periodic table covers is unknown. The problem of its upper limit is perhaps the main mystery of the periodic table. The heaviest element that has been discovered in nature is plutonium (Z = 94). The limit of artificial nuclear fusion has been reached - an element with atomic number 110. The question remains open: will it be possible to obtain elements with large atomic numbers, which ones and how many? This cannot yet be answered with any certainty.

Using complex calculations performed on electronic computers, scientists tried to determine the structure of atoms and evaluate the most important properties of “superelements,” right down to huge serial numbers (Z = 172 and even Z = 184). The results obtained were quite unexpected. For example, in an atom of an element with Z = 121, an 8p electron is expected to appear; this is after the formation of the 8s subshell has completed in atoms with Z = 119 and 120. But the appearance of p-electrons after s-electrons is observed only in atoms of elements of the second and third periods. Calculations also show that in elements of the hypothetical eighth period, the filling of the electron shells and sub-shells of atoms occurs in a very complex and unique sequence. Therefore, assessing the properties of the corresponding elements is a very difficult problem. It would seem that the eighth period should contain 50 elements (Z = 119–168), but, according to calculations, it should end at the element with Z = 164, i.e. 4 serial numbers earlier. And the “exotic” ninth period, it turns out, should consist of 8 elements. Here is his “electronic” entry: 9s 2 8p 4 9p 2. In other words, it would contain only 8 elements, like the second and third periods.

It is difficult to say how true the calculations made using a computer would be. However, if they were confirmed, then it would be necessary to seriously reconsider the patterns underlying the periodic table of elements and its structure.

The periodic table has played and continues to play a huge role in the development of various fields of natural science. It was the most important achievement of atomic-molecular science, contributed to the emergence of the modern concept of “chemical element” and clarification of concepts about simple substances and compounds.

The regularities revealed by the periodic system had a significant impact on the development of the theory of atomic structure, the discovery of isotopes, and the emergence of ideas about nuclear periodicity. The periodic system is associated with a strictly scientific formulation of the problem of forecasting in chemistry. This was manifested in the prediction of the existence and properties of unknown elements and new features of the chemical behavior of elements already discovered. Nowadays, the periodic system represents the foundation of chemistry, primarily inorganic, significantly helping to solve the problem of chemical synthesis of substances with predetermined properties, the development of new semiconductor materials, the selection of specific catalysts for various chemical processes, etc. And finally, the periodic system is the basis of teaching chemistry.