The first chemical element in the periodic table. Periodic system of elements
Periodic system chemical elements- 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 the "Experience of a system of elements based on their atomic weight and chemical affinity", compiled by D. I. Mendeleev on March 1, 1869. For two years, the scientist continuously improved the "Experience of the System", introduced the idea of groups, series and periods 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. 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 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. 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. 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 atomic nucleus of the element, and not on the atomic mass.
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 current 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 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 126 elements. Of these, all transuranium elements (Z = 93-107), 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 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 shapes(spatial and planar), analytical curves (spirals, etc.), etc. The short, long and ladder forms of tables are most widely used.
At present, preference is given to the short one.
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. Groups, in turn, are divided into main (a) and secondary (b) subgroups. Each subgroup contains elements - chemical analogues. Elements of 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 collection of elements that begins with an alkali metal and ends with an inert gas ( a special case- the first period). Each period contains a strictly defined number of elements. The periodic system consists of eight groups and eight periods.
Peculiarity first period is that it contains only 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 Iaα- or in VIIaα - subgroup, the latter option being used more often. 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. This option is not perfect, since the inertness of helium and neon is beyond doubt.
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, and in nitrogen the highest oxidation state +5 corresponds to the group number; for fluorine, the oxidation state is known to be +7. 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 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, 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.
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, 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.
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.
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.
Seventh period starts with france. Like the sixth, it must also contain 32 elements. Francium and radium, respectively, are elements of the Iaα- and IIaα-subgroups, actinium belongs to the IIIb-subgroup. 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.
Now let's see how the properties of chemical elements change in subgroups of the 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 the IIIaα-VIIaα subgroups. For metals of the Iaα-IIIaα-subgroups, an increase in chemical activity is observed. In the elements of the IVaα - VIIaα 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
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 \u003d 0 (s-elements) and ends with an element with the same value of n and l \u003d 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 ...
On the attached color tab, the symbols of the elements of each type (s-, p-, d- and f-elements) are depicted on a certain 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 the electronic configurations of the outer electron shells, which basically determine Chemical properties elements.
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 characterized by high chemical activity, because its only 1s electron is easily split off. At the same time, the configuration of helium (1s 2) is very stable, which causes its complete chemical inactivity.
For elements of a-subgroups, the outer electron shells are filled (with n equal to the period number); 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 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. 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 uncompleted 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 we saw 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.
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 RuO 4 and OsO 4 . For elements of the 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. 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, electrons 6d and 5f 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 the heptavalent 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 118. 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 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 85 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 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 to what extent the calculations made with the help of a computer would correspond to the truth. 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 of various fields of natural science. 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 oh and connections.
The laws 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. 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.
The graphic representation of the Periodic Law is the Periodic System (table). The horizontal rows of the system are called periods, and the vertical columns are called groups.
In total, there are 7 periods in the system (table), and the period number is equal to the number of electron layers in the atom of the element, the number of the external (valence) energy level, and the value of the main quantum number for the highest energy level. Each period (except the first) begins with an s-element - an active alkali metal and ends with an inert gas, which is preceded by a p-element - an active non-metal (halogen). If we move along the period from left to right, then with an increase in the charge of the nuclei of atoms of chemical elements of small periods, the number of electrons at the external energy level will increase, as a result of which the properties of the elements change - from typically metallic (because there is an active alkali metal at the beginning of the period), through amphoteric (the element exhibits the properties of both metals and non-metals) to non-metallic (active non-metal - halogen at the end of the period), i.e. metallic properties gradually weaken and non-metallic ones increase.
In large periods, with increasing nuclear charge, the filling of electrons is more difficult, which explains a more complex change in the properties of elements compared to elements of small periods. So, in even rows of long periods, with increasing nuclear charge, the number of electrons in the outer energy level remains constant and equal to 2 or 1. Therefore, while the electrons are filling the level following the outer (second from the outside), the properties of elements in even rows change slowly. In the transition to odd rows, with an increase in the charge of the nucleus, the number of electrons in the external energy level increases (from 1 to 8), the properties of the elements change in the same way as in small periods.
DEFINITION
Vertical columns in the Periodic system are groups of elements with similar electronic structure and being chemical analogues. Groups are designated by Roman numerals from I to VIII. The main (A) and secondary (B) subgroups are distinguished, the first of which contain s- and p-elements, the second - d - elements.
The subgroup number A indicates the number of electrons in the outer energy level (the number of valence electrons). For elements of B-subgroups, there is no direct relationship between the group number and the number of electrons in the outer energy level. In A-subgroups, the metallic properties of the elements increase, and the non-metallic properties decrease with increasing charge of the nucleus of the element's atom.
There is a relationship between the position of the elements in the Periodic system and the structure of their atoms:
- atoms of all elements of the same period have an equal number of energy levels, partially or completely filled with electrons;
— atoms of all elements of A subgroups have an equal number of electrons at the external energy level.
A plan for characterizing a chemical element based on its position in the Periodic Table
Usually, a characteristic of a chemical element based on its position in the Periodic system is given according to the following plan:
- indicate the symbol of the chemical element, as well as its name;
- indicate the serial number, number of the period and group (type of subgroup) in which the element is located;
- indicate the nuclear charge, mass number, number of electrons, protons and neutrons in the atom;
- write down the electronic configuration and indicate the valence electrons;
- draw electron-graphic formulas for valence electrons in the ground and excited (if possible) states;
- indicate the family of the element, as well as its type (metal or non-metal);
- compare the properties of a simple substance with the properties of simple substances formed by elements neighboring in a subgroup;
- compare the properties of a simple substance with the properties of simple substances formed by elements neighboring in a period;
- indicate the formulas of higher oxides and hydroxides with brief description their properties;
- indicate the values of the minimum and maximum oxidation states of a chemical element.
Characteristics of a chemical element using magnesium (Mg) as an example
Consider the characteristics of a chemical element using the example of magnesium (Mg) according to the plan described above:
1. Mg - magnesium.
2. Ordinal number - 12. The element is in period 3, in group II, A (main) subgroup.
3. Z=12 (nuclear charge), M=24 (mass number), e=12 (number of electrons), p=12 (number of protons), n=24-12=12 (number of neutrons).
4. 12 Mg 1s 2 2s 2 2p 6 3s 2 – electronic configuration, valence electrons 3s 2 .
5. Basic state
excited state
6. s-element, metal.
7. The highest oxide - MgO - exhibits the main properties:
MgO + H 2 SO 4 \u003d MgSO 4 + H 2 O
MgO + N 2 O 5 \u003d Mg (NO 3) 2
As a magnesium hydroxide, the base Mg (OH) 2 corresponds, which exhibits all the typical properties of bases:
Mg(OH) 2 + H 2 SO 4 = MgSO 4 + 2H 2 O
8. The degree of oxidation "+2".
9. The metallic properties of magnesium are more pronounced than those of beryllium, but weaker than those of calcium.
10. The metallic properties of magnesium are less pronounced than those of sodium, but stronger than those of aluminum (neighboring elements of the 3rd period).
Examples of problem solving
EXAMPLE 1
Exercise | Characterize the chemical element sulfur based on its position in the Periodic Table of D.I. Mendeleev |
Solution | 1. S - sulfur. 2. Ordinal number - 16. The element is in the 3rd period, in the VI group, A (main) subgroup. 3. Z=16 (nuclear charge), M=32 (mass number), e=16 (number of electrons), p=16 (number of protons), n=32-16=16 (number of neutrons). 4. 16 S 1s 2 2s 2 2p 6 3s 2 3p 4 – electronic configuration, valence electrons 3s 2 3p 4 . 5. Basic state
excited state
6. p-element, non-metal. 7. The highest oxide - SO 3 - exhibits acidic properties: SO 3 + Na 2 O \u003d Na 2 SO 4 8. The hydroxide corresponding to the higher oxide - H 2 SO 4, exhibits acidic properties: H 2 SO 4 + 2NaOH \u003d Na 2 SO 4 + 2H 2 O 9. Minimum oxidation state "-2", maximum - "+6" 10. The non-metallic properties of sulfur are less pronounced than those of oxygen, but stronger than those of selenium. 11. The non-metallic properties of sulfur are more pronounced than those of phosphorus, but weaker than those of chlorine (adjacent elements in the 3rd period). |
EXAMPLE 2
Exercise | Describe the chemical element sodium based on its position in the Periodic Table of D.I. Mendeleev |
Solution | 1. Na - sodium. 2. Ordinal number - 11. The element is in period 3, in group I, A (main) subgroup. 3. Z=11 (nuclear charge), M=23 (mass number), e=11 (number of electrons), p=11 (number of protons), n=23-11=12 (number of neutrons). 4. 11 Na 1s 2 2s 2 2p 6 3s 1 – electronic configuration, valence electrons 3s 1 . 5. Basic state 6. s-element, metal. 7. The highest oxide - Na 2 O - exhibits the main properties: Na 2 O + SO 3 \u003d Na 2 SO 4 As sodium hydroxide, the base NaOH corresponds, which exhibits all the typical properties of bases: 2NaOH + H 2 SO 4 \u003d Na 2 SO 4 + 2H 2 O 8. The oxidation state "+1". 9. The metallic properties of sodium are more pronounced than those of lithium, but weaker than those of potassium. 10. The metallic properties of sodium are more pronounced than those of magnesium (the neighboring element of the 3rd period). |
The periodic system of chemical elements is a classification of chemical elements created by D. I. Mendeleev on the basis of the periodic law discovered by him in 1869.
D. I. Mendeleev
According to the modern formulation of this law, in a continuous series of elements, arranged in order of increasing magnitude of the positive charge of the nuclei of their atoms, elements with similar properties are periodically repeated.
The periodic system of chemical elements, presented in the form of a table, consists of periods, series and groups.
At the beginning of each period (with the exception of the first) there is an element with pronounced metallic properties (alkali metal).
Symbols for the color table: 1 - chemical sign of the element; 2 - name; 3 - atomic mass (atomic weight); 4 - serial number; 5 - distribution of electrons over the layers.
As the atomic number of the element increases, equal to the value of the positive charge of the nucleus of its atom, the metallic properties gradually weaken and the non-metallic properties increase. The penultimate element in each period is an element with pronounced non-metallic properties (), and the last is an inert gas. In period I there are 2 elements, in II and III - 8 elements each, in IV and V - 18 elements each, in VI - 32 and in VII (incomplete period) - 17 elements.
The first three periods are called small periods, each of them consists of one horizontal row; the rest - in large periods, each of which (excluding the VII period) consists of two horizontal rows - even (upper) and odd (lower). In even rows of large periods are only metals. The properties of the elements in these rows change slightly with increasing serial number. The properties of elements in odd series of large periods change. In period VI, lanthanum is followed by 14 elements that are very similar in chemical properties. These elements, called lanthanides, are listed separately under the main table. Actinides, the elements following actinium, are similarly presented in the table.
The table has nine vertical groups. The group number, with rare exceptions, is equal to the highest positive valence of the elements of this group. Each group, excluding zero and eighth, is divided into subgroups. - main (located to the right) and side. In the main subgroups, with an increase in the serial number, the metallic properties of the elements are enhanced and the non-metallic properties of the elements are weakened.
Thus, chemical and series physical properties elements are determined by the place that a given element occupies in the periodic system.
Biogenic elements, i.e., elements that make up organisms and perform a certain biological role, occupy the upper part of the periodic table. The cells occupied by the elements that make up the bulk (more than 99%) of living matter are colored blue, the cells occupied by microelements are colored pink (see).
The Periodic Table of Chemical Elements is the biggest achievement modern natural science and a vivid expression of the most general dialectical laws of nature.
See also , Atomic weight.
The periodic system of chemical elements is a natural classification of chemical elements created by D. I. Mendeleev on the basis of the periodic law discovered by him in 1869.
In the original formulation, the periodic law of D. I. Mendeleev stated: the properties of chemical elements, as well as the forms and properties of their compounds, are in a periodic dependence on the magnitude of the atomic weights of the elements. Later, with the development of the doctrine of the structure of the atom, it was shown that a more accurate characteristic of each element is not the atomic weight (see), but the value of the positive charge of the nucleus of the atom of the element, equal to the ordinal (atomic) number of this element in the periodic system of D. I. Mendeleev . The number of positive charges on the nucleus of an atom is equal to the number of electrons surrounding the nucleus of an atom, since atoms as a whole are electrically neutral. In the light of these data, the periodic law is formulated as follows: the properties of chemical elements, as well as the forms and properties of their compounds, are in a periodic dependence on the positive charge of the nuclei of their atoms. This means that in a continuous series of elements arranged in ascending order of the positive charges of the nuclei of their atoms, elements with similar properties will be periodically repeated.
The tabular form of the periodic system of chemical elements is presented in its modern form. It consists of periods, series and groups. A period represents a sequential horizontal row of elements arranged in ascending order of the positive charge of the nuclei of their atoms.
At the beginning of each period (with the exception of the first) there is an element with pronounced metallic properties (alkali metal). Then, as the serial number increases, the metallic properties of the elements gradually weaken and the non-metallic properties of the elements increase. The penultimate element in each period is an element with pronounced non-metallic properties (halogen), and the last is an inert gas. Period I consists of two elements, the role of an alkali metal and a halogen is simultaneously performed by hydrogen. II and III periods include 8 elements each, called Mendeleev typical. IV and V periods have 18 elements each, VI-32. VII period is not yet completed and is replenished with artificially created elements; there are currently 17 elements in this period. I, II and III periods are called small, each of them consists of one horizontal row, IV-VII - large: they (with the exception of VII) include two horizontal rows - even (upper) and odd (lower). In even rows of large periods, only metals are found, and the change in the properties of the elements in the row from left to right is weakly expressed.
In odd series of large periods, the properties of the elements in the series change in the same way as the properties of typical elements. In an even number of the VI period after lanthanum 14 elements follow [called lanthanides (see), lanthanides, rare earth elements], similar in chemical properties to lanthanum and to each other. Their list is given separately under the table.
Separately, the elements following the actinium-actinides (actinides) are written out and given under the table.
There are nine vertical groups in the periodic table of chemical elements. The group number is equal to the highest positive valency (see) of the elements of this group. The exceptions are fluorine (it happens only negatively monovalent) and bromine (it does not happen heptavalent); in addition, copper, silver, gold can exhibit a valence greater than +1 (Cu-1 and 2, Ag and Au-1 and 3), and of the elements of group VIII, only osmium and ruthenium have a valence of +8. Each group, with the exception of the eighth and zero, is divided into two subgroups: the main (located to the right) and the secondary. The main subgroups include typical elements and elements of large periods, the secondary - only elements of large periods and, moreover, metals.
In terms of chemical properties, the elements of each subgroup of this group differ significantly from each other, and only the highest positive valency is the same for all elements of this group. In the main subgroups, from top to bottom, the metallic properties of elements increase and non-metallic ones weaken (for example, francium is an element with the most pronounced metallic properties, and fluorine is non-metallic). Thus, the place of an element in the periodic system of Mendeleev (serial number) determines its properties, which are the average of the properties of neighboring elements vertically and horizontally.
Some groups of elements have special names. So, the elements of the main subgroups of group I are called alkali metals, group II - alkaline earth metals, group VII - halogens, elements located behind uranium - transuranium. Elements that are part of organisms, take part in metabolic processes and have a pronounced biological role, are called biogenic elements. All of them occupy the upper part of the table of D. I. Mendeleev. This is primarily O, C, H, N, Ca, P, K, S, Na, Cl, Mg and Fe, which make up the bulk of living matter (more than 99%). The places occupied by these elements in the periodic table are colored in light blue. Biogenic elements, which are very few in the body (from 10 -3 to 10 -14%), are called microelements (see). In the cells of the periodic system, stained in yellow, trace elements are placed, the vital importance of which for humans has been proven.
According to the theory of the structure of atoms (see Atom), the chemical properties of elements depend mainly on the number of electrons in the outer electron shell. Periodic change in the properties of elements with an increase in positive charge atomic nuclei due to the periodic repetition of the structure of the outer electron shell (energy level) of atoms.
In small periods, with an increase in the positive charge of the nucleus, the number of electrons in the outer shell increases from 1 to 2 in period I and from 1 to 8 in periods II and III. Hence the change in the properties of the elements in the period from an alkali metal to an inert gas. The outer electron shell containing 8 electrons is complete and energetically stable (elements of the zero group are chemically inert).
In large periods in even rows, with an increase in the positive charge of the nuclei, the number of electrons in the outer shell remains constant (1 or 2) and the second outer shell is filled with electrons. Hence the slow change in the properties of elements in even rows. In odd series of long periods, with an increase in the charge of the nuclei, the outer shell is filled with electrons (from 1 to 8) and the properties of the elements change in the same way as for typical elements.
The number of electron shells in an atom is equal to the period number. The atoms of the elements of the main subgroups have a number of electrons on their outer shells equal to the group number. The atoms of the elements of the secondary subgroups contain one or two electrons on the outer shells. This explains the difference in the properties of the elements of the main and secondary subgroups. The group number indicates the possible number of electrons that can participate in the formation of chemical (valence) bonds (see Molecule), therefore such electrons are called valence. For elements of secondary subgroups, not only the electrons of the outer shells, but also the penultimate ones, are valence. The number and structure of electron shells are indicated in the attached periodic table of chemical elements.
The periodic law of D. I. Mendeleev and the system based on it have exclusively great importance in science and practice. The periodic law and the system were the basis for the discovery of new chemical elements, the precise determination of their atomic weights, the development of the theory of the structure of atoms, the establishment of geochemical laws for the distribution of elements in earth's crust and development contemporary ideas about living matter, the composition of which and the laws associated with it are in accordance with the periodic system. The biological activity of the elements and their content in the body are also largely determined by the place they occupy in the periodic system of Mendeleev. So, with an increase in the serial number in a number of groups, the toxicity of elements increases and their content in the body decreases. The periodic law is a vivid expression of the most general dialectical laws of the development of nature.
In nature, there are a lot of repeating sequences:
- seasons;
- Times of Day;
- days of the week…
In the middle of the 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 miraculous dreams of the scientist was the Periodic Table of Chemical Elements, in which D.I. Mendeleev arranged the chemical elements in order of increasing atomic mass. In the modern table, the chemical elements are arranged in ascending order of the atomic number of the element (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). Note that the atomic mass of some elements is a non-integer! Remember isotopes! Atomic mass is the weighted average of all isotopes of an element that occur naturally under natural conditions.
Below the table are the lanthanides and actinides.
Metals, non-metals, metalloids
They are located in the Periodic Table to the left of the stepped diagonal line that starts 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 Periodic table. Basic properties of metals: solid (except mercury); glitter; good electrical and thermal conductors; plastic; malleable; donate electrons easily.
The elements to the right of the stepped diagonal B-Po are called non-metals. The properties of non-metals are directly opposite to the properties of metals: poor conductors of heat and electricity; fragile; non-forged; non-plastic; usually accept electrons.
Metalloids
Between metals and non-metals are semimetals(metalloids). They are characterized by the properties of both metals and non-metals. Semimetals have found their main industrial application in the production of semiconductors, without which no modern microcircuit or microprocessor is inconceivable.
Periods and groups
As mentioned above, the periodic table consists of seven periods. In each period, the atomic numbers of the elements increase from left to right.
The properties of elements in periods change sequentially: so sodium (Na) and magnesium (Mg), which are 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 the IA(1) group, 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 been given special names. These groups are discussed below.
Group IA(1). The atoms of the elements of this group have only one electron in the outer electron layer, so they easily donate one electron.
The most important alkali metals are sodium (Na) and potassium (K) as they play important role in the process of human life and are part of the 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). The atoms of the elements of this group have two electrons in the outer electron layer, which 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 the elements of this group usually receive one electron each, because. on the outer electronic layer there are five elements each, and one electron is just missing to the "complete set".
The most famous elements of this group are: 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 the elements of this group have a fully "staffed" outer electron layer. Therefore, they "do not 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 Barlett discovered that some of these gases, under certain conditions, can still react with other elements.
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 see 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 each:
- 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 Valence and the Table of electronic configurations of atoms of chemical elements by periods.
Let us now turn our attention to the elements located in groups with symbols AT. They are located in the center periodic table and are called transition metals.
A distinctive feature of these elements is the presence of electrons in atoms 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
Separate from the main table are located lanthanides and actinides 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 chem. elements, their natures. , which is a table expression . The prototype of the periodic. chemical systems. elements was the table "Experience of a system of elements based on their and chemical similarity", compiled by D. I. Mendeleev on March 1, 1869 (Fig. 1). In the last For years, the scientist improved the table, developed ideas about periods and groups of elements and about the place of the element in the system. In 1870 Mendeleev called the system natural, and in 1871 periodic. As a result, even then the periodic system largely acquired the modern. structural outlines. Based on it, Mendeleev predicted the existence and St. Islands approx. 10 unknown elements; these predictions were subsequently confirmed.
Rice. 1 Table "Experience of a system of elements based on their and chemical similarity" (D. I. Mendeleev. I myrtle, 1869).
However, over the next more than 40 years, the periodic system means. degree was only empirical. generalization of facts, since there was no physical. explanation of the causes of periodic changes in CB-B elements depending on the increase in their . Such an explanation was impossible without reasonable ideas about the structure (see). Therefore, the most important milestone in the development of the periodic system 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 system is numerically equal to posit. 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 a strict physical. formulation, it was possible to unambiguously determine the lower. boundary of the periodic system (H as an element with min. 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 system was developed in the beginning. 1920s (see below).
The structure of the periodic system. 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 nature 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 symbols corresponding to them in the tables; for an element with Z = 109, the maximum values are still unknown. long-lived.
Over the entire history of the periodic system, more than 500 different versions of its image have been published. This was due to attempts to find a rational solution to some controversial problems of the structure of the periodic system (location of H, lanthanides, etc.). Naib. spread following. tabular forms of the expression of the periodic system: 1) a short one was proposed by Mendeleev (in its modern form it is placed at the beginning of the volume on a 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, the short and long forms have been especially widely used as visual and practically convenient. All listed. forms have certain advantages and disadvantages. However, it is hardly possible to offer k.-l. universal a variant of the image of the periodic system, to-ry would adequately reflect the whole variety of St. in chem. elements and the specifics of changes in their chemical. behavior as Z increases.
Fundam. the principle of constructing the periodic system 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 \u003d 118) and 8 groups. a collection of elements beginning (or the first period) and ending with . The number of elements in periods naturally increases and, starting from the second, they 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 value 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 similar in chem. St. you, to-rykh are characterized by the same structure of external. electronic shells. In most groups, elements of subgroups a and b show a certain chem. similarity, prim. in higher.
Special place group VIII occupies in the structure of the periodic system. Throughout the duration time, only the elements of "triads" were attributed to it: Fe-Co-Ni and (Ru Rh Pd and Os-Ir-Pt), and all were placed in their own. zero group; therefore, the periodic system contained 9 groups. After in the 60s. Comm. were received. Xe, Kr and Rn began to be placed in subgroup VIIIa, and the zero group was abolished. The elements of the triads constituted subgroup VIII6. Such a "structural design" of group VIII now appears in almost all published versions of the expression of the periodic system.
Distinguish. The feature of the first period is that it contains only 2 elements: H and He. due to St-in - unities. an element that does not have a well-defined place in the periodic table. The symbol H is placed either in subgroup Ia, or in subgroup VIIa, or both at the same time, enclosing the symbol in brackets in one of the subgroups, or, finally, depicting it decomp. fonts. These ways of arranging H are based on the fact that it has some formal similarities with both with and with .
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 (ones, + 1); followed by Be(+2). metallic character B (+3) is weakly expressed, and the character C following it is typical (+4). The subsequent N, O, F and Ne-non-metals, and only in N the highest + 5 corresponds 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 change in chem. st-in to-rykh is in many respects similar to that observed in the second period. However, Mg and Al are more "metallic" than resp. Be and B. The remaining elements are Si, P, S, Cl and Ar are non-metals; they all exhibit , equal to the group number, except for 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. the 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, 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 alkali-earth. Ca (s-elements) follows a series of 10 so-called. transitional (Sc-Zn), or d-elements (symbols of blue color), which are included in subgroups b. Most (all of them - ) exhibit higher values 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 st-in is in many ways similar to the change in st-in of elements of the second and third periods in the corresponding intervals of Z values. For Kr, several. relatively stable Comm., in DOS. with F.
The fifth period (Rb-Xe) is constructed similarly to the fourth; it also has an insert of 10 transitional, or d-elements (Y-Cd). Features of changes in St-in elements in the period: 1) in the triad Ru-Rh-Pd shows max, 4-8; 2) all elements of subgroups a, including Xe, exhibit higher values equal to the group number; 3) I have weak metallic. sv. T. arr., the properties of the elements of the fourth and fifth periods as Z increases are more difficult to change than the properties of the elements in the second and third periods, which is primarily due to the presence of transitional 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 chem. St. to you (preferably in +3) and therefore not m. b. placed in different system groups. In the short form of the periodic system, all lanthanides are included in subgroup IIIa (La), and their totality is deciphered under the table. This technique is not without drawbacks, since 14 elements seem 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 triad Os Ir Pt, only Os exhibits max. +8; 2) At is more pronounced compared to I metallic. character; 3) Rn max. reactive from , but strong makes it difficult to study its chem. sv.
The seventh period, like the sixth, should contain 32 elements, but is not yet completed. Fr and Ra elements resp. subgroups Ia and IIa, Ac analogue of the elements of subgroup III6. According to the actinide concept of G. Seaborg (1944), Ac is followed by a family of 14 f-elements (Z = 90 103). In the short form of the periodic system, the latter are included in Ac and are similarly written as otd. line below the table. This technique assumed the presence of a certain chem. similarities of 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, stabilization of the lower ones is typical for heavy ones (+2 or even +1 for Md).
Assessment of chem. the nature of Ku (Z = 104) and Ns (Z = 105), synthesized in the number of single very short-lived ones, led to the conclusion that these elements are analogues, respectively. Hf and Ta, i.e., d-elements, and should be placed in subgroups IV6 and V6. Chem. elements with Z= 106 109 have not been studied, 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 (subgroups IIIa VIIIa).
Theory of the Periodic System was premier. created by H. (1913 21) on the basis of the quantum model he proposed. Taking into account the specifics of the change in the properties of elements in the periodic system and information about them, he developed a scheme for constructing electronic configurations as Z increases, using it as 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 are periodically repeated, which determines the periodicity. change in chem. sv-in elements. This is ch. cause of physical the nature of the phenomenon of periodicity. Electronic shells, with the exception of those that correspond to the values 1 and 2 of the main quantum number l, are not filled sequentially and monotonously until they are completely completed (the numbers in successive shells are: 2, 8, 18, 32, 50, ... ); their construction is periodically interrupted by the appearance of collections (constituting certain subshells), which correspond to large values of n. This is the essence. feature of the "electronic" interpretation of the structure of the periodic system.
The scheme for the formation of electronic configurations, which underlies the theory of the periodic system, reflects, i.e., a certain sequence of appearance in as Z grows, sets (subshells) characterized by certain values of the principal and orbital (l) quantum numbers. This scheme in general view is written in the form of a table. (see below).
Vertical lines separate subshells, which are filled in the elements that make up the sequence. periods of the periodic system (numbers of periods are indicated by numbers at the top); subshells that complete the formation of shells with the given item are highlighted in bold.
The numbers in shells and subshells are defined by . With regard to , as particles with a half-integer , he postulates that in not m. two with the same values of all quantum numbers. The capacitances of 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 2p |
3s 3p |
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 a 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 is filled with the same n and l = 1 (np 6 -elements) you); the exception is the first period (only 1s elements). All s- and p-elements belong to subgroups a. Subgroups b include elements in which shells that previously remained unfinished are being completed (the values of h are less than the period number, l = 2 and 3). The first three periods include elements of only subgroups a, i.e., s- and p-elements.
The real 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 subsequent increase in the sum (n + /). In this case, 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 does not appear in La with Z = 57, but in the next Ce (Z = 58); follow. the construction of the 4f subshell is interrupted in Gd (Z = 64, the presence of a 5d electron). Such a "blurring of periodicity" clearly affects the seventh period for with Z > 89, which is reflected in the properties of the elements.
The real scheme was not originally derived from the c.-l. strict theoretical. representations. It was based on well-known chem. Holy Islands of elements and information about their spectra. Valid. physical substantiation of the real scheme was due to the application of methods to the description of the structure. In quantum mech. interpretation of the theory of structure, the concept of electron 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, in the first approximation, quite exhaustively explains the theoretical. bases of the periodic system. In any case, the published forms of the representation of the periodic system reflect the idea of the nature of the distribution over shells and subshells.
Structure and chemical properties of elements. Main features of chem. the behavior of the 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 1s-elements of the first period (H and He). due to the presence in only one, a largesv. The configuration of He (1s 2) is exceptional, which determines its chem. inertia. Since the elements of the subgroups a are filled with ext. electron shells (with n equal to the number of the period), St-va elements change markedly as Z increases in the corresponding periods, which is expressed in the weakening of the metal and the strengthening of the non-metallic. sv. All but H and He are p-elements. At the same time, in each subgroup a, as Z increases, an increase in metallicity is observed. sv. These patterns are explained by the weakening of the binding energy of ext. with the kernel during the transition from period to period.
The value of the periodic system. This system has played and continues to play a huge role in the development of many. natural science. disciplines. She became important link in atomic mol. teachings, contributed to the formulation of modern. the concept of "chemical element" and the clarification of ideas about simple in-wah and Comm., rendered means. influence on the development of the theory of structure and the emergence of the concept of isotopy. With the periodic system is connected strictly scientific. statement of the forecasting problem in , whichmanifested itself both in the prediction of the existence of unknown elements and their properties, and new features of the chemical. behavior of already exposed elements. The periodic system is the most important basis of inorg. ; it serves, for example, tasks synthesis in-in with predetermined St. you, the creation of new materials, in particular semiconductor, the selection of specific. for diff. chem. processes. Periodic system - scientific. base of teaching general and non-org. , as well as certain branches of atomic physics.
Lit .: Mendeleev D.I., Periodic law. Main articles, M., 1958; Kedrov B. M.. Three aspects of atomistics, part 3. Mendeleev's law, M., 1969; Trifonov D H., 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. Mixed chronology major events. Moscow, 1974; Karapetyami MX. Drakii S. I., Structure, M., 1978; The doctrine of periodicity. History and modernity. Sat. articles. M.. 1981. Korolkov D.V., Osnovy, 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 displacement is called the vector connecting the start and end points of the trajectory The vector connecting the beginning and end of the path is called
- Trajectory, path length, displacement vector Vector connecting the initial position
- Calculating the area of a polygon from the coordinates of its vertices The area of a triangle from the coordinates of the vertices formula
- Acceptable Value Range (ODZ), theory, examples, solutions