Variants of the periodic system of chemical elements. Lecture on the topic: "The Periodic Table of Chemical Elements D.I.

Periodic Table of Elements D. I. Mendeleev, natural, which is a tabular (or other graphical) expression. The periodic system of elements was developed by D. I. Mendeleev in 1869-1871.

Story periodic system elements. Attempts to systematize were made by various scientists in England, the USA from the 30s of the 19th century. Mendeleev - I. Döbereiner, J. Dumas, French chemist A. Shancourtua, eng. chemists W. Odling, J. Newlands, and others established the existence of groups of elements that are similar in chemical properties, the so-called "natural groups" (for example, Döbereiner's "triad"). However, these scientists did not go further than establishing particular patterns within groups. In 1864, L. Meyer proposed a table showing the ratio for several characteristic groups of elements on the basis of data on. Meyer did not make theoretical reports from his table.

The prototype of the scientific periodic system of elements was the table "Experience of a system of elements based on their and chemical similarity", compiled by Mendeleev on March 1, 1869 ( rice. one). Over the next two years, the author improved this table, introduced ideas about groups, series and periods of elements; made an attempt to estimate the capacity of small and large periods, containing, in his opinion, 7 and 17 elements, respectively. In 1870 he called his system natural, and in 1871 - periodic. Even then, the structure of the periodic system of elements acquired in many respects modern outlines ( rice. 2).

The periodic system of elements did not immediately win recognition as a fundamental scientific generalization; the situation changed significantly only after the discovery of Ga, Sc, Ge and the establishment of the divalence of Be (it was considered trivalent for a long time). Nevertheless, the periodic system of elements largely represented an empirical generalization of facts, since the physical meaning of the periodic law was unclear and there was no explanation for the reasons for the periodic change in the properties of elements depending on the increase. Therefore, up to the physical substantiation of the periodic law and the development of the theory of the periodic system of elements, many facts could not be explained. So, unexpected was the discovery at the end of the 19th century. , which seemed to have no place in the periodic table of elements; this difficulty was eliminated due to the inclusion in the periodic system of elements of an independent zero group (subsequently VIIIa-subgroup). The discovery of many "radio elements" at the beginning of the 20th century. led to a contradiction between the need for their placement in the periodic system of elements and its structure (for more than 30 such elements there were 7 "vacant" places in the sixth and seventh periods). This contradiction was overcome as a result of the discovery. Finally, the value () as a parameter that determines the properties of elements gradually lost its meaning.

One of the main reasons for the impossibility of explaining the physical meaning of the periodic law and the periodic system of elements was the lack of a theory of structure (see, Atomic physics). Therefore, the most important milestone in the development of the periodic system of elements was the planetary model proposed by E. Rutherford (1911). On its basis, the Dutch scientist A. van den Broek suggested (1913) that an element in the periodic system of elements (Z) is numerically equal to the charge of the nucleus (in units of elementary charge). This was experimentally confirmed by G. Moseley (1913-14, see Moseley's law). So it was possible to establish that the frequency of changing the properties of elements depends on , and not on . As a result, on a scientific basis, the lower limit of the periodic system of elements was determined (as an element with a minimum Z = 1); the number of elements between and is exactly estimated; it was found that the "gaps" in the periodic system of elements correspond to unknown elements with Z = 43, 61, 72, 75, 85, 87.

However, the question of the exact number remained unclear, and (which is especially important) the reasons for the periodic change in the properties of elements depending on Z were not revealed. These reasons were found in the course of further development of the theory of the periodic system of elements based on quantum ideas about the structure (see. Further). The physical substantiation of the periodic law and the discovery of the phenomenon of isotonia made it possible to scientifically define the concept of "" (""). The attached periodic table (see. ill.) contains modern meanings elements on the carbon scale in accordance with the International Table 1973. The most long-lived are given in square brackets. Instead of the most stable 99 Tc, 226 Ra, 231 Pa and 237 Np, these are indicated, adopted (1969) by the International Commission on.

Structure of the Periodic Table of the Elements. The modern (1975) Periodic Table of Elements covers 106; of these, all transuranium (Z = 93-106), as well as elements with Z = 43 (Tc), 61 (Pm), 85 (At) and 87 (Fr) were obtained artificially. Throughout the history of the periodic system of elements, it has been proposed a large number of(several hundred) variants of its graphic representation, mainly in the form of tables; known images and in the form of various geometric shapes(spatial and planar), analytical curves (for example, ), etc. The most widely used are three forms of the periodic system of elements: the short one proposed by Mendeleev ( rice. 2) and received universal recognition (in its modern form, it is given on ill.); long ( rice. 3); staircase ( rice. 4). The long form was also developed by Mendeleev, and in an improved form it was proposed in 1905 by A. Werner. The ladder form was proposed by the English scientist T. Bailey (1882) and the Danish scientist J. Thomsen (1895) and improved by N. (1921). Each of three forms has advantages and disadvantages. The fundamental principle of building a periodic system of elements is the division of all into groups and periods. Each group, in turn, is divided into the main (a) and secondary (b) subgroups. Each subgroup contains elements that have similar chemical properties. The elements of the a- and b-subgroups in each group, as a rule, show a certain chemical similarity among themselves, mainly in the higher ones, which, as a rule, correspond to the group number. A period is a collection of elements that begins and ends (a special case is the first period); Each period contains a strictly defined number of elements. The Periodic Table of the Elements consists of 8 groups and 7 periods (the seventh has not yet been completed).

The specificity of the first period is that it contains only 2 elements: H and He. The place of H in the system is ambiguous: since it exhibits properties common with and with, it is placed either in the Ia- or (more preferably) in the VIIa-subgroup. - the first representative of the VIIa-subgroup (however, for a long time He and all were united into an independent zero group).

The second period (Li - Ne) contains 8 elements. It starts with Li, the only one of which is I. Then comes Be - , II. The metallic nature of the next element B is weakly expressed (III). The C following it is typical, it can be both positively and negatively tetravalent. Subsequent N, O, F and Ne - , and only in N the highest V corresponds to the group number; only in rare cases shows positive, and VI is known for F. Ends period Ne.

The third period (Na - Ar) also contains 8 elements, the nature of the change in the properties of which is largely similar to that observed in the second period. However, Mg, unlike Be, is more metallic, as is Al compared to B, although Al is inherent. Si, P, S, Cl, Ar are typical, but all of them (except Ar) show higher values ​​equal to the group number. Thus, in both periods, as Z increases, a weakening of the metallic and strengthening of the non-metallic nature of the elements is observed. Mendeleev called the elements of the second and third periods (small, in his terminology) typical. It is significant that they are among the most common in nature, and C, N and O, along with H, are the main elements of organic matter (organogens). All elements of the first three periods are included in subgroups a.

According to modern terminology (see below), the elements of these periods refer to s-elements (alkaline and alkaline earth) that make up the Ia- and IIa-subgroups (highlighted in red on the color table), and p-elements (B - Ne, At - Ar), included in IIIa - VIIIa-subgroups (their symbols are highlighted orange). For elements of small periods, with increasing, first, a decrease is observed, and then, when the number in the outer shell already increases significantly, their mutual repulsion leads to an increase in . The next maximum is reached at the beginning of the next period on an alkaline element. Approximately the same pattern is typical for.

The fourth period (K - Kr) contains 18 elements (the first large period, according to Mendeleev). After K and alkaline earth Ca (s-elements), there follows a series of ten so-called (Sc - Zn), or d-elements (symbols are given in blue), which are included in subgroups of 6 corresponding groups of the periodic table of elements. Most (all of them) show higher ones equal to the group number. An exception is the triad Fe - Co - Ni, where the last two elements are maximally positively trivalent, and under certain conditions it is known in VI. Elements starting with Ga and ending with Kr (p-elements) belong to subgroups a, and the nature of the change in their properties is the same as in the corresponding intervals Z for elements of the second and third periods. It has been established that Kr is able to form (mainly with F), but VIII is unknown for it.

The fifth period (Rb - Xe) is constructed similarly to the fourth; it also has an insert of 10 (Y - Cd), d-elements. Specific features of the period: 1) in the triad Ru - Rh - Pd only shows VIII; 2) all elements of subgroups a exhibit higher values ​​equal to the group number, including Xe; 3) I has weak metallic properties. Thus, the nature of the change in properties with increasing Z for elements of the fourth and fifth periods is more complicated, since the metallic properties are preserved in large interval.

The sixth period (Cs - Rn) includes 32 elements. In addition to 10 d-elements (La, Hf - Hg), it contains a set of 14 f-elements, from Ce to Lu (black symbols). The elements La to Lu are chemically very similar. In the short form of the periodic table, the elements are included in La (because their predominant III) and are written on a separate line at the bottom of the table. This technique is somewhat inconvenient, since 14 elements are, as it were, outside the table. The long and ladder forms of the Periodic Table of Elements do not have such a shortcoming, reflecting well the specifics against the background of the integral structure of the Periodic Table of Elements. Features of the period: 1) in the triad Os - Ir - Pt only manifests VIII; 2) At has a more pronounced (compared to 1) metallic character; 3) Rn, apparently (it is little studied), should be the most reactive of .

The seventh period, starting from Fr (Z = 87), should also contain 32 elements, of which 20 are known so far (before the element with Z = 106). Fr and Ra are elements of the Ia- and IIa-subgroups (s-elements), respectively, Ac is an analogue of the elements of the IIIb subgroup (d-element). The next 14 elements, the f-elements (with Z from 90 to 103), make up the . In the short form of the periodic system of elements, they occupy Ac and are written in a separate line at the bottom of the table, like , in contrast to which they are characterized by a significant variety. In connection with this, in the chemical respect, the series and reveal noticeable differences. Study of chemical nature elements with Z = 104 and Z = 105 showed that these elements are analogues of and, respectively, that is, d-elements, and should be placed in IVb- and Vb-subgroups. Subsequent elements up to Z = 112 should also be members of b-subgroups, and then (Z = 113-118) p-elements (IIIa - VIlla-subgroups) will appear.

Theory of the periodic system of elements. The theory of the periodic system of elements is based on the idea of ​​the specific patterns of construction of electron shells (layers, levels) and subshells (shells, sublevels) in as Z increases (see, Atomic physics). This idea was developed in 1913-21, taking into account the nature of the change in properties in the periodic system of elements and the results of studying them. revealed three significant features of the formation of electronic configurations: 1) the filling of electron shells (except for shells corresponding to the values ​​of the main quantum number n = 1 and 2) does not occur monotonously until their full capacity, but is interrupted by the appearance of sets related to shells with large values n; 2) similar types of electronic configurations are periodically repeated; 3) the boundaries of the periods of the periodic system of elements (with the exception of the first and second ones) do not coincide with the boundaries of successive electron shells.

In the notation adopted in atomic physics, the real scheme of the formation of electronic configurations as Z increases can be general view written as follows:

Vertical lines separate the periods of the periodic system of elements (their numbers are indicated by numbers at the top); subshells are highlighted in bold, which complete the construction of shells with a given n. Under the designations of subshells, the values ​​of the main (n) and orbital (l) quantum numbers are given, which characterize successively filled subshells. In accordance with the capacity of each electron shell is 2n 2, and the capacity of each subshell is 2(2l + 1). From the above scheme, the capacities of successive periods are easily determined: 2, 8, 8, 18, 18, 32, 32... Each period begins with an element in which it appears with a new value of n. Thus, periods can be characterized as collections of elements starting with an element with the value n equal to the period number and l = 0 (ns 1 -elements), and ending with an element with the same n and l = 1 (np 6 -elements); the exception is the first period containing only ls elements. In this case, the a-subgroups include elements for which n is equal to the period number, and l \u003d 0 or 1, that is, an electron shell is built with a given n. The b-subgroups include elements in which the completion of shells that remained unfinished (in this case n is less than the period number, and l = 2 or 3). The first - third periods of the periodic system of elements contain only elements of a-subgroups.

The given real scheme for the formation of electronic configurations is not perfect, since in some cases the clear boundaries between successively filled subshells are violated (for example, after filling in Cs and Ba 6s subshells, not a 4f-, but a 5d-electron appears in Cs and Ba, there is a 5d-electron in Gd etc.). Moreover, the original circuit could not have been deduced from any fundamental physical concepts; such a conclusion was made possible by its application to the problem of structure.

Types of configurations of external electron shells (on ill. configurations are indicated) determine the main features of the chemical behavior of the elements. These features are specific for elements of a-subgroups (s- and p-elements), b-subgroups (d-elements) and f-families ( and ). A special case are the elements of the first period (H and He). The high chemical atomic number is explained by the ease of splitting off a single ls-electron, while the configuration (1s 2) is very strong, which causes it chemical inertness.

Since the elements of a-subgroups fill the outer electron shells (with n equal to the period number), the properties of the elements change noticeably as Z increases. So, in the second period Li (configuration 2s 1) is chemically active, easily losing valence, a Be (2s 2) - also, but less active. The metallic nature of the next element B (2s 2 p) is weakly expressed, and all subsequent elements of the second period, in which the 2p subshell is built up, are narrower. The eight-electron configuration of the outer electron shell Ne (2s 2 p 6) is extremely strong, therefore - . A similar character of the change in properties is observed for the elements of the third period and for s- and p-elements all subsequent periods, however, the weakening of the connection between the outer and the core in a-subgroups as Z increases in a certain way affects their properties. So, in s-elements, there is a noticeable increase in chemical properties, and in p-elements, an increase in metallic properties. In the VIIIa subgroup, the stability of the ns 2 np 6 configuration is weakened, as a result of which already Kr (the fourth period) acquires the ability to enter into. The specificity of the p-elements of the 4th-6th periods is also connected with the fact that they are separated from the s-elements by sets of elements in which the previous electron shells are built up.

For transitional d-elements of b-subgroups, incomplete shells are completed with n one less than the period number. The configuration of their outer shells, as a rule, is ns 2 . Therefore, all d-elements are . The similar structure of the outer shell of the d-elements in each period leads to the fact that the change in the properties of the d-elements as Z increases is not sharp and a clear difference is found only in higher ones, in which the d-elements show a certain similarity with the p-elements of the corresponding groups of the periodic element systems. The specificity of the elements of the VIIIb subgroup is explained by the fact that their d-subshells are close to completion, and therefore these elements (with the exception of Ru and Os) are not inclined to exhibit higher . For elements of the Ib subgroup (Cu, Ag, Au), the d-subshell actually turns out to be complete, but not yet sufficiently stabilized; these elements also show higher ones (up to III in the case of Au).

The value of the periodic system of elements. The periodic system of elements has played and continues to play a huge role in the development of natural science. It was the most important achievement of the atomic and molecular theory, made it possible to give a modern definition of the concept "" and clarify the concepts of and compounds. The patterns revealed by the periodic system of elements had a significant impact on the development of the theory of structure, contributed to the explanation of the phenomenon of isotony. A strictly scientific formulation of the problem of forecasting in is associated with the periodic system of elements, which manifested itself both in predicting the existence of unknown elements and their properties, and in predicting new features of the chemical behavior of already discovered elements. Periodic system of elements - foundation, primarily inorganic; it significantly helps to solve problems of synthesis with predetermined properties, develop new materials, in particular semiconductor materials, and select materials specific for various chemical processes etc. Periodic system of elements - also scientific basis teaching .

Lit .: Mendeleev D.I., Periodic law. Main articles, M., 1958; Kedrov BM, Three aspects of atomistics. Part 3. Zakon Mendeleev, M., 1969; Rabinovich E., Tilo E., Periodic system of elements. History and theory, M.-L., 1933; Karapetyants M. Kh., Drakin S. I., Structure, M., 1967; Astakhov K. V., Current state periodic system of D. I. Mendeleev, M., 1969; Kedrov B. M., Trifonov D. N., The law of periodicity and. Discoveries and chronology, M., 1969; One Hundred Years of the Periodic Law. Collection of articles, M., 1969; One Hundred Years of the Periodic Law. Reports at plenary sessions, M., 1971; Spronsen J. W. van, The periodic system of chemical elements. A history of the first hundred years, Amst.-L.-N.Y., 1969; Klechkovsky V. M., The distribution of atomic and the rule of successive filling of (n + l)-groups, M., 1968; Trifonov D. N., On the quantitative interpretation of periodicity, M., 1971; Nekrasov B.V., Fundamentals, vol. 1-2, 3rd ed., M., 1973; Kedrov B. M., Trifonov D. N., On modern problems of the periodic system, M., 1974.

D. N. Trifonov.

Rice. 1. Table "Experience of a system of elements", based on their and chemical similarity, compiled by D. I. Mendeleev on March 1, 1869.

Rice. 3. Long form of the periodic system of elements (modern version).

Rice. 4. Ladder form of the periodic system of elements (according to N., 1921).

Rice. 2. "Natural system of elements" by D. I. Mendeleev (short form), published in the 2nd part of the 1st edition of the Fundamentals in 1871.

Periodic system of elements of D. I. Mendeleev.

Periodic system - an ordered set of chemical elements, their natural classification, which is a graphical (tabular) expression of the periodic law of chemical elements. Its structure, in many respects 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 “Experiment of a system of elements based on their atomic weight and chemical similarity”, compiled by D. I. Mendeleev on March 1, 1869. For two and a half years, the scientist continuously improved the “Experience of the System”, introduced the idea of ​​​​groups, series and periods of elements. As a result, the structure of the periodic system acquired in many respects modern outlines.

Important for its evolution was the concept of the place of an element in the system, determined by the numbers of the group and period. Based on this concept, Mendeleev came to the conclusion that it is necessary to change the atomic masses of some elements: uranium, indium, cerium and its satellites. It was the first practical use periodic system. Mendeleev was also the first to predict the existence and properties of several unknown elements. The scientist described the most important properties ekaaluminum (the future gallium), ekabora (scandium) and ekasilicium (germanium). In addition, he predicted the existence of analogues of manganese (future technetium and rhenium), tellurium (polonium), iodine (astatine), cesium (francium), 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 system.

The first versions of the periodic system in many respects represented only an empirical generalization. After all, the physical meaning of the periodic law was not clear, there was no explanation of the reasons for the periodic change in the properties of elements depending on the increase in atomic masses. As a result, many problems remained unresolved. Are there limits to the periodic system? Is it possible to determine the exact number of existing elements? The structure of the sixth period remained unclear - what is the exact amount of rare earth elements? It was not known whether there are still elements between hydrogen and lithium, what is the structure of the first period. 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. Unexpected was the discovery in 1894-1898. 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 system. 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 system and its existing structure. For them, there were only 7 vacancies in the sixth and seventh periods. This problem was solved as a result of 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 not known how the atom was arranged (see Atom). The most important milestone in the development of the periodic system 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 ordinal number of an element in the periodic system 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 Z - the charge of the nucleus of an atom of an element, and not on atomic mass (see Periodic law of chemical elements).

As a result, the structure of the periodic system has been significantly strengthened. The lower bound of the system has been determined. This is hydrogen, the element with the minimum Z = 1. It has become possible to accurately estimate the number of elements between hydrogen and uranium. "Gaps" in the periodic system 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 periodic change in the properties of elements were not revealed. depending on Z.

Based on the established structure of the periodic system and the results of the study of atomic spectra, the Danish scientist N. Bohr in 1918–1921. developed ideas about the sequence of construction of electron 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 electron shells and subshells of atoms.

The periodic system 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 (francium) 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, as well as 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, the short form is preferred.

The fundamental principle of building the periodic system is its division into groups and periods. Mendeleev's concept of rows of elements is not currently used, since it is devoid of physical meaning. The groups, in turn, are subdivided into the main (a) and secondary (b) subgroups. Each subgroup contains elements - chemical analogues. The elements of the a- and b-subgroups in most groups also show a certain similarity among themselves, mainly in higher oxidation states, which, as a rule, are equal to the group number. A period is a set 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 system consists of eight groups and seven periods, and the seventh period has not yet been completed.

Peculiarity first period lies in the fact that it contains only 2 gaseous elements in the free form: hydrogen and helium. The place of hydrogen in the system is ambiguous. Since it exhibits properties in common with alkali metals and halogens, it is placed either in the 1a- or Vlla-subgroup, or 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 provision required revision after the synthesis chemical compounds krypton, xenon and radon. As a result, inert gases and elements of the former group VIII (iron, cobalt, nickel and platinum metals) were combined into one group.

Second 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 the boron, carbon is a typical non-metal 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 non-metals. The inert gas neon completes 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 the elements of the second period. But there is also its own specificity. So, magnesium, unlike beryllium, is more metallic, as well as aluminum compared to boron. Silicon, phosphorus, sulfur, chlorine, argon are all typical non-metals. And all of them, except for argon, exhibit the highest oxidation states equal to the group number.

As we can see, in both periods, as Z increases, a distinct weakening of the metallic and strengthening of the non-metallic properties of the elements is observed. D. I. Mendeleev called the elements of the second and third periods (in his words, small ones) typical. The elements of small periods are among the most common in nature. Carbon, nitrogen and oxygen (along with hydrogen) are organogens, that is, 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, a series of elements follows, 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 for iron, cobalt, and nickel. Elements from gallium to krypton belong to the a-subgroups. A number of chemical compounds are known for krypton.

Fifth period (rubidium - xenon) in its construction is similar 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 the a‑subgroups exhibit the highest oxidation states equal to the group number. The features of the change in the properties of the elements of the fourth and fifth periods as Z grows are more complex in comparison with the second and third periods.

Sixth period (cesium - radon) includes 32 elements. In this period, in addition to 10 transition metals (lanthanum, hafnium - mercury), there is also a set of 14 lanthanides - from cerium to lutetium. The 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 system, the lanthanide series 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 is the most reactive of all inert gases. Unfortunately, due to the fact that it is highly radioactive, its chemistry has been little studied (see Radioactive Elements).

Seventh period starts with france. Like the sixth, it should also contain 32 elements, but 24 of them are known so far. Francium and radium, respectively, are elements of subgroups Ia and IIa, actinium belongs to subgroup IIIb. Next comes the actinide family, which includes elements from thorium to lawrencium and is arranged similarly to the lanthanides. The decoding of this row 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 nature of the elements as Z increases. This pattern is especially pronounced in IIIa–VIIa subgroups. For metals of Ia–IIIa‑subgroups, an increase in chemical activity is observed. In the elements of IVa–VIIa‑subgroups, as Z increases, a weakening of the chemical activity of the elements is observed. For elements of b‑subgroups, 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 1920s. 20th 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 system occurs in the following sequence:

Based on the theory of the periodic system, the following definition of a period can be given: a period is a collection of elements that begins with an element with a value of n equal to the period number and l = 0 (s-elements) and ends with an element with the same value of n and l = 1 (p- 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 shown on a specific color background: s-elements - on red, p-elements - on orange, d-elements - on blue, f-elements - on green. In each cell, the serial numbers and atomic masses of the elements are given, as well as electronic configurations outer electron shells.

It follows from the theory of the periodic system that elements with n equal to the period number and l = 0 and 1 belong to the a-subgroups. The b-subgroups include those elements in whose atoms the shells that previously remained incomplete are completed. That is why the first, second and third periods do not contain elements of b‑subgroups.

The structure of the periodic system 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 are periodically repeated. Namely, they determine the main features of the chemical behavior of elements. These features manifest themselves differently for the elements of the a-subgroups (s- and p-elements), for the elements of the b-subgroups (transitional d-elements) and the 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 highly reactive because its only 1s electron is easily split off. At the same time, the configuration of helium (1s 2) is very stable, which makes it chemically inactive.

For elements of a-subgroups, the outer electron shells of 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 (configuration 2s) is an active metal that easily loses a single valence electron ; beryllium (2s 2) is also a metal, but less active due to the fact that its outer electrons are more firmly bound to the nucleus. Further, boron (2s 2 p) has a weakly pronounced metallic character, and all subsequent elements of the second period, in which the construction of a 2p subshell occurs, are already nonmetals. 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 the elements of the second period are explained by the desire of their atoms to acquire the electronic configuration of the nearest inert gas (the helium configuration for elements from lithium to carbon or the neon configuration for elements from carbon to fluorine). This is why, for example, oxygen cannot exhibit a higher oxidation state equal to the group number: after all, it is easier for it to achieve the neon configuration by acquiring additional electrons. The same nature of the change in properties is manifested 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 the outer electrons and the nucleus in a-subgroups as Z increases manifests itself in the properties of the corresponding elements. So, for s-elements, there is a noticeable increase in chemical activity as Z increases, and for p-elements, an increase in metallic properties.

In atoms of transitional d-elements, previously unfinished shells are completed with the value of the main quantum number n, one less than the period number. With some exceptions, the configuration of the outer electron shells of transition element atoms 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 sharp as is observed in s- and p-elements. In higher oxidation states, d-elements show a certain similarity with p-elements of the corresponding groups of the periodic system.

Features 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, are not inclined to give compounds of higher oxidation states. The only exceptions are ruthenium and osmium, which give the oxides RuO 4 and OsO 4 . For elements of Ib- and IIb-subgroups, the d-subshell actually turns out to be complete. Therefore, they exhibit oxidation states equal to the group number.

In the atoms of lanthanides and actinides (all of them are metals), the completion of previously incomplete electron shells occurs with the value of the main quantum number n 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 outside N shell is filled with 4f electrons. That's why the lanthanides are so similar.

For actinides, the situation is more complicated. In atoms of elements with Z = 90–95, electrons 6d and 5f 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 act in the heptavalent state. Only elements starting from curium (Z = 96) become stable in the trivalent state, but even here there are some peculiarities. Thus, the properties of the actinides differ significantly from those of the lanthanides, and therefore both families cannot be considered similar.

The actinide family ends with an element with Z = 103 (lawrencium). An evaluation 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 family of actinides in atoms, the systematic filling of the 6d subshell begins. The chemical nature of elements with Z = 106–110 has not been experimentally evaluated.

The finite number of elements that the periodic system covers is unknown. The problem of its upper limit is, perhaps, the main riddle of the periodic system. The heaviest element found in nature is plutonium (Z = 94). The reached limit of artificial nuclear fusion is an element with the atomic number 110. The question remains: will it be possible to obtain elements with higher atomic numbers, which ones and how many? It cannot yet be answered with any certainty.

Using the most complex calculations performed on electronic computers, scientists tried to determine the structure of atoms and evaluate the most important properties of "superelements", up 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, the appearance of an 8p electron is expected; this is after the formation of the 8s subshell was completed in the 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 the elements of the hypothetical eighth period, the filling of the electron shells and sub-shells of atoms occurs in a very complex and peculiar sequence. Therefore, to evaluate 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., by 4 sequence numbers before. And the "exotic" ninth period, it turns out, should consist of 8 elements. Here is his "electronic" record: 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 with the help of a computer would be. However, if they were confirmed, then it would be necessary to seriously revise the patterns underlying the periodic system of elements and its structure.

The periodic system has played and continues to play a huge role in the development various areas natural sciences. It was the most important achievement of atomic and molecular science, contributed to the emergence modern concept"chemical element" and clarifying the concepts of simple substances and compounds.

The patterns revealed by the periodic system had a significant impact on the development of the theory of the structure of atoms, the discovery of isotopes, and the emergence of ideas about nuclear periodicity. A strictly scientific statement of the problem of forecasting in chemistry is connected with the periodic system. This manifested itself in the prediction of the existence and properties of unknown elements and new features of the chemical behavior of elements already discovered. Now the periodic system is 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 underlies teaching chemistry.

How to use the periodic table? For an uninitiated person, reading the periodic table is the same as looking at the ancient runes of elves for a dwarf. And the periodic table, by the way, if used correctly, can tell a lot about the world. In addition to serving you in the exam, it is also simply indispensable in solving problems. huge amount chemical and physical problems. But how to read it? Fortunately, today everyone can learn this art. In this article we will tell you how to understand the periodic table.

The periodic system of chemical elements (Mendeleev's table) is a classification of chemical elements that establishes the dependence of various properties of elements on the charge of the atomic nucleus.

History of the creation of the Table

Dmitri Ivanovich Mendeleev was not a simple chemist, if someone thinks so. He was a chemist, physicist, geologist, metrologist, ecologist, economist, oilman, aeronaut, instrument maker and teacher. During his life, the scientist managed to conduct a lot of fundamental research in the most different areas knowledge. For example, it is widely believed that it was Mendeleev who calculated the ideal strength of vodka - 40 degrees. We do not know how Mendeleev treated vodka, but it is known for sure that his dissertation on the topic “Discourse on the combination of alcohol with water” had nothing to do with vodka and considered alcohol concentrations from 70 degrees. With all the merits of the scientist, the discovery of the periodic law of chemical elements - one of the fundamental laws of nature, brought him the widest fame.

There is a legend according to which the scientist dreamed of the periodic system, after which he only had to finalize the idea that had appeared. But if it were that easy... This version about the creation of the periodic table, apparently, is nothing more than a legend. When asked how the table was opened, Dmitry Ivanovich himself answered: “ I’ve been thinking about it for maybe twenty years, and you think: I sat and suddenly ... it’s ready. ”

In the middle of the nineteenth century, attempts to streamline the known chemical elements (63 elements were known) were simultaneously undertaken by several scientists. For example, in 1862 Alexandre Emile Chancourtois placed elements along a helix and marked cyclic repetition chemical properties. Chemist and musician John Alexander Newlands proposed his version of the periodic table in 1866. An interesting fact is that in the arrangement of the elements the scientist tried to discover some mystical musical harmony. Among other attempts was the attempt of Mendeleev, which was crowned with success.

In 1869, the first scheme of the table was published, and the day of March 1, 1869 is considered the day of the discovery of the periodic law. The essence of Mendeleev's discovery was that the properties of elements with increasing atomic mass do not change monotonously, but periodically. The first version of the table contained only 63 elements, but Mendeleev undertook a number of very non-standard solutions. So, he guessed to leave a place in the table for yet undiscovered elements, and also changed the atomic masses of some elements. The fundamental correctness of the law derived by Mendeleev was confirmed very soon after the discovery of gallium, scandium and germanium, the existence of which was predicted by scientists.

Modern view of the periodic table

Below is the table itself.

Today, instead of atomic weight (atomic mass), the concept of atomic number (the number of protons in the nucleus) is used to order elements. The table contains 120 elements, which are arranged from left to right in ascending order of atomic number (number of protons)

The columns of the table are so-called groups, and the rows are periods. There are 18 groups and 8 periods in the table.

• The metallic properties of elements decrease when moving along the period from left to right, and in reverse direction- increase.
• The dimensions of atoms decrease as they move from left to right along the periods.
• When moving from top to bottom in the group, the reducing metallic properties increase.
• Oxidizing and non-metallic properties increase along the period from left to right. I am.

What do we learn about the element from the table? For example, let's take the third element in the table - lithium, and consider it in detail.

First of all, we see the symbol of the element itself and its name under it. In the upper left corner is the atomic number of the element, in the order in which the element is located in the table. The atomic number, as already mentioned, is equal to the number of protons in the nucleus. The number of positive protons is usually equal to the number of negative electrons in an atom (with the exception of isotopes).

The atomic mass is indicated under the atomic number (in this version of the table). If we round the atomic mass to the nearest integer, we get the so-called mass number. The difference between the mass number and the atomic number gives the number of neutrons in the nucleus. Thus, the number of neutrons in a helium nucleus is two, and in lithium - four.

So our course "Mendeleev's Table for Dummies" has ended. In conclusion, we suggest you watch a thematic video, and we hope that the question of how to use periodic table th Mendeleev, became more understandable to you. Reminder to study new item always more effective not alone, but with the help of an experienced mentor. That is why, you should never forget about those who will gladly share their knowledge and experience with you.

The graphical representation of the periodic law is Periodic system chemical elements. More than \(700\) forms of the periodic table are known. official by decision International Union chemists is its semi-long version.

Each chemical element in the table is assigned one cell, in which the symbol and name of the element, serial number and relative atomic mass are indicated.

The broken line marks the boundary between metals and non-metals.

The sequence of arrangement of elements does not always coincide with the increase in atomic mass. There are several exceptions to the rule. So, the relative atomic mass of argon is less than the atomic mass of potassium, in tellurium it is less than that of iodine.

Each element has its own ordinal (atomic) room , is located in a certain period and a certain group.

A period is a horizontal row of chemical elements starting with an alkali metal (or hydrogen) and ending with an inert (noble) gas.

Table seven periods. Each contains a certain number of elements:

\(1\)th period - \(2\) elements,

\(2\)th period - \(8\) elements,

\(3\)th period - \(8\) elements,

\(4\)th period - \(18\) elements,

\(5\)th period - \(18\) elements,

\(6\)th period - \(32\) element (\(18 + 14\)),

The \(7\)th period is the \(32\) element (\(18 + 14\)).

The first three periods are called small periods, the rest big . In both small and large periods, there is a gradual weakening of metal properties and reinforcement of non-metallic , only in large periods it occurs more smoothly.

Elements with serial numbers \(58\)–\(71\) ( lanthanides ) and \(90\)–\(103\) ( actinides ) are removed from the table and placed below it. These are group IIIB elements. The lanthanides belong to sixth period, and actinides - to seventh .

The eighth period will appear in the Periodic Table when new elements are discovered.

Group - a vertical column of chemical elements that have similar properties.

In the Periodic Table, \(18\) groups are numbered with Arabic numerals. Roman numerals are often used with the addition of the letters \(A\) or \(B\). In this case, the groups \(8\).

Groups \(A\) begin with elements of small periods, also include elements of large periods; contain both metals and non-metals. In the short version of the Periodic Table, this is main subgroups .

The groups \(B\) contain elements of large periods, and these are only metals. In the short version of the Periodic Table, this is side subgroups .

Number of elements in groups:

IA , VIIIA - by \(7\) elements;

IIA - VIIA - by \(6\) elements;

IIIB - \ (32 \) element (\ (4 + 14 \) lanthanides \ (+ 14 \) actinides);

VIIIB - \(12\) elements;

IB , IIB , IVB - VIIB - by \(4\) element.

The number of groups will change as new elements are added to the table.

Roman group number usually indicates higher valency in oxides. But for some elements this rule is not fulfilled. So, fluorine is not heptavalent, but oxygen - hexavalent. Do not show a valence equal to the group number, helium , neon and argon These elements do not form compounds with oxygen. Copper is divalent and gold - trivalent, although these are elements of the first group.

Lecture # 2

Periodic system of chemical elements D.I. Mendeleev

Plan:

Discovery of D.I. Mendeleev Periodic Law

The principle of constructing the periodic system

The periodic law in the formulation of D.I. Mendeleev.

The periodic system of chemical elements is a natural classification of chemical elements, which is a graphical (tabular) expression of the periodic law of chemical elements. Its structure, in many respects 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 " Experience of a system of elements based on their atomic weight and chemical similarity", compiled DI. Mendeleev March 1, 1869. Over the course of two years, the scientist continuously improved the "Experience of the System", introduced the concept of groups, series and periods of elements. As a result, the structure of the periodic system acquired in many respects modern outlines.

Important for its evolution was the concept of the place of an element in the system, determined by the numbers of the group and period. Based on this concept, Mendeleev came to the conclusion that it is necessary to change the atomic masses of some chemical elements: uranium, indium, cerium and its satellites. This was the first practical application of the periodic system. Mendeleev was also the first to predict the existence of several unknown elements. The scientist described the most important properties of ekaaluminum (future gallium), ekabor (scandium) and ekasilicon (germanium). In addition, he predicted the existence of analogues of manganese (future technetium and rhenium), tellurium (polonium), iodine (astatine), cesium (francium), 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 system.

The first versions of the periodic system of chemical elements in many respects represented only an empirical generalization. After all, the physical meaning of the periodic law was not clear, there was no explanation of the reasons for the periodic change in the properties of elements depending on the increase in atomic masses. As a result, many problems remained unresolved. Are there limits to the periodic system? Is it possible to determine the exact number of existing elements? The structure of the sixth period remained unclear - what is the exact amount of rare earth elements. It was not known whether there are still elements between hydrogen and lithium, what is the structure of the first period. 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. Unexpected was the discovery in 1894 - 1898. a galaxy of 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 system. 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 system and its existing structure. For them, there were only 7 vacancies in the sixth and seventh periods. This problem was solved as a result of the establishment of shift rules and the discovery of isotopes.

One of the main reasons for the inability to explain the physical meaning of the periodic law and the structure of the periodic system was that it was not known how the atom was built (see Atom). The most important milestone in the development of the periodic system 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 ordinal number of an element in the periodic system 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 an atom of an element, and not on atomic mass.

As a result, the structure of Mendeleev's periodic system was significantly strengthened. The lower bound of the system has been determined. This is hydrogen, the element with the minimum Z = 1. It has become possible to accurately estimate the number of elements between hydrogen and uranium. "Gaps" in the periodic system 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 periodic change in the properties of elements were not revealed. depending on Z.

Based on the current structure of the periodic system and the results of the study of atomic spectra, the Danish scientistN. Bor in 1918 - 1921gg. developed ideas about the sequence of construction of electron shells and subshells in atoms. The scientist came to the conclusion that similar types of electronic configurations 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 electron shells and subshells of atoms.

Currently, the periodic system covers 117 elements.Of these, all transuranium elements (Z "= 93 - 117), as well as elements with Z = 43 (technetium), 61 (promethium), 85 (astatine), 87 (francium) were obtained artificially. Over the entire history of the existence of the periodic system, it was proposed a large number (> 500) of variants of its graphic representation, mainly in the form of tables, as well as in the form of various geometric figures (spatial and planar), analytical curves (spirals, etc.), etc. The most widely used are short, long and stair forms of periodic tables.Currently, preference is given to a short one.

fundamental principle construction of the periodic system is itsdivision into groups and periods.Mendeleev's concept of rows of elements is not currently used, since it is devoid of physical meaning.Groups, in turn, are divided into main (a) and secondary (b) subgroups.Each subgroup contains elements - chemical analogues. The elements of the a- and b-subgroups in most groups also show a certain similarity, mainly in higher oxidation states, which, as a rule, are equal to the group number.

A period is a set 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 system consists of eight groups and seven periods, and the seventh is not yet completed.

The peculiarity of the first period isin that it containsonly 2 elements: hydrogen and helium. The place of hydrogen in the system is ambiguous. Since it exhibits properties in common with alkali metals and halogens, it is placed either in the I A- or in the VII A-subgroup, the latter option being used more often. Helium is the first representative of the VIII A-subgroup. For a long time, helium and all inert gases were separated into an independent zero group. This provision required revision after the synthesis of chemical compounds of krypton, xenon and radon. As a result, inert gases and elements of the former group VIII (iron, cobalt, nickel and platinum metals) were combined into one group. This option is not perfect, since the inertness of helium and neon is beyond doubt.

The second period contains 8 elements.It begins with the alkali metal lithium, whose only oxidation state is +1. This is followed by beryllium (a metal, oxidation state +2). Boron already exhibits a weakly pronounced metallic character and is a non-metal (oxidation state + 3). The next carbon is a typical non-metal, which 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; for fluorine, the oxidation state + 7 is known. The inert gas neon completes the period.

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

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

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

The 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, a series of elements follows, consisting of 10 so-called transition metals (scandium - zinc). All of them belong to b-subgroups. Most transition metals exhibit higher oxidation states equal to the group number, except for iron, cobalt, and nickel. Elements from gallium to krypton belong to the A-subgroups. Krypton, unlike the previous inert gases, can form chemical compounds.

The fifth period (rubidium - xenon) is similar in construction 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 the A-subgroups exhibit the highest oxidation states equal to the group number, excluding xenon. It can be seen that the features of the change in the properties of the elements of the fourth and fifth periods as Z grows are more complex in comparison with the second and third periods.

The sixth period (cesium - radon) includes 32 elements.In this period, in addition to 10 transition metals (lanthanum, hafnium - mercury), there is also a set of 14 lanthanides - from cerium to lutetium. The 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 system, a number of lanthanum species are included in the lanthanum cell and the decoding of this series is given at the bottom of the table.

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 is probably the most reactive of all inert gases. Unfortunately, due to the fact that it is highly radioactive, its chemistry has been little studied.)

The seventh period begins with France.Like the sixth, it should also contain 32 elements, but 21 of them are known so far. Francium and radium, respectively, are elements of I a- and I I a-subgroups, actinium belongs to III b-subgroup. The further construction of the seventh period is debatable. The most common idea is about the actinide family, which includes elements from thorium to lawrencium and is similar to the lanthanides. The decoding of this row of elements is also given at the bottom of the table.

How do the properties of chemical elements change in the subgroups of the periodic system of Mendeleev?

The main pattern of this change is the strengthening of the metallic nature of the elements as Z increases. This pattern is especially pronounced in the IIIa-VIIa subgroups. For metals I A-III A-subgroups, an increase in chemical activity is observed. In the elements of IVA - VIIA subgroups, as Z increases, a weakening of the chemical activity of the elements is observed. For elements of b-subgroups, the change in chemical activity is more difficult.

The theory of the periodic system was developed by N. Bohr and other scientists in the 1920s.20th century and is based on a real scheme for the formation of electronic configurations of atoms. 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 system occurs in the following sequence:

Period numbers

1 2 3 4 5 6 7

1s2s2p 3s3p 4s3d4p 5s4d5p 6s4f5d6p7s5f6d7p

Based on the theory of the periodic system, the following definition of a period can be given:a period is a collection of elements beginning with the element with 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). The exception is the first period containing only 1s elements. The number of elements in periods follows from the theory of the periodic system: 2, 8, 8. 18, 18, 32 ...

The b-subgroups include those elements in whose atoms 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 system of chemical elements is closely related to the structure of atoms of chemical elements. As Z increases, similar types of configuration of the outer electron shells are periodically repeated. Namely, they determine the main features of the chemical behavior of elements. These features manifest themselves differently for the elements of the A-subgroups (s- and p-elements), for the elements of the b-subgroups (transitional d-elements), and for the 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 b-electron is easily split off. At the same time, the configuration of helium (1st) is very stable, which causes its complete chemical inactivity.

The elements of the A-subgroups are filled with outer electron shells (with n equal to the number of the period); therefore, the properties of these elements change markedly as Z increases. Thus, in the second period, lithium (configuration 2s) is an active metal that easily loses its only valence electron; beryllium (2s~) is also a metal, but less active due to the fact that its outer electrons are more firmly bound to the nucleus. Further, boron (2s "p) has a weakly pronounced 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 ~ p ~) - an inert gas - is very durable.

The chemical properties of the elements of the second period are explained by the desire of their atoms to acquire the electronic configuration of the nearest inert gas (the helium configuration for elements from lithium to carbon or the neon configuration for elements from carbon to fluorine). This is why, for example, oxygen cannot exhibit a higher oxidation state equal to the group number: after all, it is easier for it to achieve the neon configuration by acquiring additional electrons. The same nature of the change in properties is manifested 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 the outer electrons and the nucleus in A-subgroups as Z increases manifests itself 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, an increase in metallic properties.

In atoms of transitional d-elements, previously unfinished shells with the value of the main quantum number and one less than the period number are completed. With some exceptions, the configuration of the outer electron shells of transition element atoms is ns. Therefore, all d-elements are metals, and that is why the changes in the properties of 1-elements as Z increases are not as sharp as we saw with s and p-elements. In higher oxidation states, d-elements show a certain similarity with p-elements of the corresponding groups of the periodic system.

The features of the properties of the elements of triads (VIII b-subgroup) are explained by the fact that the d-subshells are close to completion. This is why iron, cobalt, nickel and platinum metals, as a rule, are not inclined to give compounds of higher oxidation states. The only exceptions are ruthenium and osmium, which give the oxides RuO4 and OsO4. For elements of I- and II B-subgroups, the d-subshell actually turns out to be complete. Therefore, they exhibit oxidation states equal to the group number.

In the atoms of lanthanides and actinides (they are all metals)there is a completion of previously unfinished electron shells with the value of the main quantum number and two units less than the period number. In the atoms of these elements, the configuration of the outer electron shell (ns2) remains unchanged. At the same time, f-electrons do not actually affect the chemical properties. That's why the lanthanides are so similar.

For actinides, the situation is much more complicated.In the range of nuclear charges Z = 90 - 95, the electrons 6d and 5/ can take part in chemical interactions. And from this it follows that actinides exhibit a much wider range of oxidation states. For example, for neptunium, plutonium and americium, compounds are known where these elements act in a seven-valence state. Only for elements starting from curium (Z = 96) does the trivalent state become stable. Thus, the properties of the actinides differ significantly from those of the lanthanides, and therefore both families cannot be considered similar.

The actinide family ends with an element with Z = 103 (lawrencium). An evaluation 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 family of actinides in atoms, the systematic filling of the 6d subshell begins.

The finite number of elements that the periodic system covers is unknown. The problem of its upper limit is, perhaps, the main riddle of the periodic system. The heaviest element found in nature is plutonium (Z = 94). The reached limit of artificial nuclear fusion is an element with the atomic number 107. The question remains: will it be possible to obtain elements with higher atomic numbers, which ones and how many? It cannot yet be answered with any certainty.

With the help of the most complex calculations performed on a computer, scientists tried to determine the structure of atoms and evaluate the most important properties of such "superelements", up to huge serial numbers (Z = 172 and even Z = 184). The results obtained were quite unexpected. For example, in the atom of an element with Z = 121, the appearance of an 8p electron is assumed; this is after the formation of the 8s subshell was completed in the 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 the elements of the hypothetical eighth period, the filling of electron shells and subshells of atoms occurs in a very complex and peculiar sequence. Therefore, to evaluate 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" record: 9s "Зр 9р". 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 with the help of a computer would be. However, if they were confirmed, then the regularities underlying the periodic system of elements and its structure would have to be seriously reconsidered.

The periodic system 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 and molecular science, contributed to the emergence of the modern concept of "chemical element" and the refinement of the concepts of simple substances and compounds.

Patterns revealed by the periodic system,had a significant impact on the development of the theory of the structure of atoms, the discovery of isotopes, the emergence of ideas about nuclear periodicity. A strictly scientific statement of the problem of forecasting in chemistry is connected with the periodic system. This manifested itself 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 is 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. Finally, the periodic system underlies the teaching of chemistry.

Periodic law of Mendeleev

The periodic law of chemical elements is a fundamental law of nature that reflects the periodic change in the properties of chemical elements as the charges of the nuclei of their atoms increase. Opened on March 1 (February 17 according to the old style) 1869 D.I. Mendeleev. On this day, he compiled a table called "The experience of a system of elements based on their atomic weight and chemical similarity." The final formulation of the periodic law was given by Mendeleev in July 1871. It read:

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

Mendeleev's formulation of the periodic law existed in science for over 40 years. It was revised thanks to the outstanding achievements of physics, mainly the development of the nuclear model of the atom. It turned out,charge of the nucleus of an atom (Z) numerically equalsserial numberof the corresponding element in the periodic system, and the filling of the electron shells and subshells of atoms, depending on Z, occurs in such a way that similar electronic configurations of atoms are periodically repeated (see Periodic system of chemical elements). Therefore, the modern formulation of the periodic law is as follows:the properties of elements, simple substances and their compounds are in a periodic dependence on the charges of the nuclei of atoms.

Unlike other fundamental laws of nature, such as the law of universal gravitation or the law of equivalence of mass and energy, the periodic law cannot be written in the form of any general equation or formula. Its visual reflection is the periodic table of elements. However, both Mendeleev himself and other scientists made attempts to find a mathematical equation for the periodic law of chemical elements. These attempts were crowned with success only after the development of the theory of the structure of the atom. But they concern only the establishment of a quantitative dependence of the order of distribution of electrons in shells and subshells on the charges of atomic nuclei.

The periodic law is a universal law for the entire universe.It is valid wherever atoms exist. But not only the electronic structures of atoms change periodically. The structure and properties of atomic nuclei also obey a peculiar periodic law. In nuclei consisting of neutrons and protons, there are neutron and proton shells, the filling of which has a periodic character. There are even attempts to construct a periodic system of atomic nuclei.

Dmitry Ivanovich Mendeleev (1834 - 1907)

Russian scientist, discovered the periodic law of chemical elements.

In 1955 Americanphysicists led by G. Seaborg synthesized a chemical element with a serial number101. They gave it a namemendelevium- in recognition of the merits of the great Russian scientist.For more than 100 years, Mendeleev's periodic system has served as the key to the discovery of new elements.

The periodic law and the periodic system became the most important contribution of D. I. Mendeleev to the development of natural science. But they are only a part of the great creative heritage of the scientist.The complete collection of his works - 25 voluminous volumes, a real encyclopedia of knowledge.

Mendeleev brought scattered information about isomorphism into a system, and this played a role in the development of geochemistry. He discovered the critical boiling point, above which a substance cannot be in a liquid state, developed the hydrate theory of solutions, and thus is rightfully considered an outstanding physical chemist. He conducted in-depth studies of the properties of rarefied gases, showing himself to be an outstanding experimental physicist. Mendeleev proposed the theory of the inorganic origin of oil, which still has adherents; developed a process for making smokeless powder; studied aeronautics, meteorology, improved the technique of measurements. Being the manager of the Main Chamber of Weights and Measures, he did a lot for the development of metrology. For his scientific merits, Mendeleev was elected a member of more than 50 academies and scientific societies around the world. V scientific activity the scientist saw, in his words, his "first service to the motherland."

Second service - pedagogical activity. Mendeleev was the author of the textbook "Fundamentals of Chemistry", which during his lifetime went through 8 editions and was translated more than once into foreign languages. Mendeleev taught in many educational institutions Petersburg. “Of the thousands of my students, there are now many prominent figures everywhere, and, meeting them, I always heard that they planted a good seed in them, and not simply served a duty,” the scientist wrote in his declining years.

The “third service to the Motherland” was multifaceted and useful - in the field of industry and Agriculture. Here Mendeleev showed himself to be a true patriot who cared about the development and future of Russia. In his estate Boblovo, he set up "experiments on the cultivation of bread." He studied in detail the methods of oil production and gave a lot valuable advice for their improvement. He constantly delved into the urgent needs of industry, visited factories and factories, mines and mines. The authority of Mendeleev was so high that he was constantly invited as an expert to solve complex problems. economic problems. Shortly before his death, he published the book "To the Knowledge of Russia", in which he outlined an extensive program for the development of the country's productive forces.

"Scientific sowing will sprout for the harvest of the people" - this was the motto of all the activities of the scientist.

Mendeleev was one of the most cultured people of his time. He was deeply interested in literature and art, collected a huge collection of reproductions of paintings by artists from different countries and peoples. Meetings of prominent cultural figures often took place at his apartment.

Control questions:

In what year was the periodic law of chemical elements discovered, as formulated by D. I. Mendeleev?

What is the essence of the law of periodicity? What are its main features?

What is a period, group, subgroup in the periodic system?

Which subgroups are called main and which are secondary?

How do the metallic properties of elements change in a group and in a period?

How do the redox properties of atoms of elements change with increasing atomic number?

In which groups of the periodic table are the elements that form gaseous compounds with hydrogen? Which of them are acidic?

If you draw a line in the periodic system from boron to astatine, then the elements with what properties will be on the left side of this line?

What is the essence of the quantum mechanical theory of the structure of the atom?

Give the modern formulation of the periodic law of D. I. Mendeleev?

Find in the periodic table an element located in the IV period, in the V row and showing valency VI in the oxygen compound. What is its hydrogen valency?

Literature:

Gabrielyan O.S. Chemistry for professions and specialties technical profile: textbook / O.S. Gabrielyan, I.G. Ostroumov. - M .: Publishing Center "Academy", 2009. - 256 p.

Gabrielyan O.S. Chemistry: studies for students. avg. prof. textbook institutions / O.S. Gabrielyan, I.G. Ostroumov. - 6th ed., Sr. - M .: Publishing Center "Academy", 2009. - 336s.