In 1871 Mendeleev's periodic law was formulated. By this time, 63 elements were known to science, and Dmitry Ivanovich Mendeleev ordered them on the basis of relative atomic mass. The modern periodic table has expanded significantly.

History

In 1869, while working on a chemistry textbook, Dmitry Mendeleev faced the problem of systematizing the material accumulated over many years by various scientists - his predecessors and contemporaries. Even before Mendeleev's work, attempts were made to systematize the elements, which served as prerequisites for the development of the periodic system.

Figure: 1. Mendeleev D.I ..

Element classification searches are summarized in the table.

Mendeleev ordered the elements according to their relative atomic mass, arranging them in ascending order. There are nineteen horizontal and six vertical rows in total. This was the first edition of the periodic table of the elements. This is where the history of the discovery of the periodic law begins.

It took the scientist almost three years to create a new, more perfect table. Six columns of elements turned into horizontal periods, each of which began with an alkali metal and ended with a non-metal (inert gases were not yet known). The horizontal rows formed eight vertical groups.

Unlike his colleagues, Mendeleev used two criteria for the distribution of elements:

  • atomic mass;
  • chemical properties.

It turned out that there is a pattern between these two criteria. After a certain number of elements with increasing atomic mass, the properties begin to repeat themselves.

Figure: 2. The table compiled by Mendeleev.

Initially, the theory was not expressed mathematically and could not be fully confirmed experimentally. The physical meaning of the law became clear only after the creation of the atomic model. The idea is to repeat the structure of the electron shells with a sequential increase in the charges of the nuclei, which is reflected in the chemical and physical propertiesah elements.

Law

Having established the periodicity of changes in properties with an increase in atomic mass, Mendeleev in 1871 formulated the periodic law, which became fundamental in chemical science.

Dmitry Ivanovich determined that the properties of simple substances are periodically dependent on the relative atomic masses.

Science of the XIX century did not have modern knowledge about the elements, therefore the modern formulation of the law is somewhat different from Mendeleev's. However, the essence remains the same.

With the further development of science, the structure of the atom was studied, which influenced the formulation of the periodic law. According to the modern periodic law, the properties chemical elements depend on charges atomic nuclei.

Table

Since the time of Mendeleev, the table he created has significantly changed and began to reflect almost all the functions and characteristics of the elements. The ability to use the table is essential for further study of chemistry. The modern table is presented in three forms:

  • short - periods occupy two lines, and hydrogen is often referred to as group 7;
  • long - isotopes and radioactive elements are removed from the table;
  • extra long - each period occupies a separate line.

Figure: 3. Long modern table.

The short table is the most obsolete version and was canceled in 1989, but is still used in many textbooks. The long and extra long shapes are internationally recognized and used all over the world. Despite the established forms, scientists continue to improve the periodic system, offering the latest options.

What have we learned?

Periodic law and Mendeleev's periodic system were formulated in 1871. Mendeleev identified the regularities of the properties of elements and ordered them on the basis of the relative atomic mass. With increasing masses, the properties of the elements changed and then repeated. Subsequently, the table was supplemented, and the law was adjusted in accordance with modern knowledge.

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The periodic law, one of the fundamental laws of natural science, was discovered by the great Russian scientist D.I. Mendeleev in 1869. Initially, the law was formulated as follows: the properties of elements and their compounds are periodically dependent on the value of their atomic weight (according to modern concepts - atomic mass).

The periodic law was presented as a classification of the elements. On its basis, the elements were arranged in natural groups according to the totality of their properties. Special attention was paid to this moment: guided by the properties of the elements, D.I. In a number of cases, Mendeleev even had to deviate from the sequential arrangement of elements in the Periodic Table strictly in terms of increasing atomic masses (atomic "weights"), for example, 18 Ar (39.9) and 19 K (39.1), 52 Te (127.6 ) and 53 1 (126.9).

At the time of Mendeleev, the reason for the periodicity of the properties of elements was not known. However, the discoverer of the Periodic Law was sure that the reason should be sought in the structure of the substance.

The discovery of the Periodic Law not only provided the foundation for chemical science, but also posed the task of elucidating the physical cause of periodicity. Chemical and the absolute majority of physical properties of elements are a periodic function of some independent, uniquely determined quantity inherent in each element and changing monotonically from element to element. Atomic mass ("atomic weight") was taken by Mendeleev as such a value.

Only when, thanks to the successes of physics, it became known much more about the structure of the atom than at the time of the discovery and formation of the periodic law, did its true meaning and reasons for periodicity become clear. From element to element according to the Periodic Table, the charge of the nucleus of the element's atom changes, which is determined by the number of protons. In the Periodic Table, this number coincides with the ordinal number of the element. Since the atom is electrically neutral, the nuclear charge (in units of the electron charge) is equal to the number of electrons in the electron shell of the atom. An increase in the ordinal number of an element by one means that one proton has been added to the nucleus of the atom, and one electron, respectively, in the electron shell. Since the properties of elements, especially chemical ones, are determined mainly by the electrons of the outer quantum layer, the reason for the periodicity of the properties is the periodicity of the nature of filling the space around the nucleus with electrons. The factor that determines the structure of the electron shells of atoms, and, consequently, the properties of elements, is the charge of the atomic nucleus. Therefore, the modern formulation of the periodic law is as follows: the properties of elements and their compounds are periodically dependent on the charge of the nucleus of the element's atom.

The atomic mass of an element is determined by the total number of nucleons (protons and neutrons) in the nuclei of the isotopes of this element and the isotopic composition of the element. The change in atomic mass is mainly proportional to the nuclear charge. Therefore, Mendeleev's formulation of the Periodic Law, with a few exceptions, correctly reflects the arrangement of the elements in the Periodic Table, but does not reveal the reason for the periodicity.

According to Pauli's principle, the number of possible electronic states in quantum levels and sublevels is limited by the number of combinations of non-repeating sets of four quantum numbers p, /, t and s, and this determines the capacity of quantum levels and sublevels (see Table 2.1). If the atom is not excited, electrons fill such orbitals, the energy on which is minimal.

The periodic table would be simpler if the main quantum number determined the energy in many-electron atoms, as in the hydrogen atom. Then, in accordance with the capacity of the quantum layers, the periods would consist of 2, 8, 18, 32, 50, etc. elements, and noble gases with a complete quantum level would have numbers 2, 10, 28, 60, 110 ... However, due to the electron-electron interaction, this sequence is violated. From the IV period, the filling of a new quantum layer, which in the Periodic system corresponds to the beginning of a new period, begins with an incomplete preliminary III quantum level, and from the VI period - with incomplete IV and V quantum levels, etc. Therefore, noble gases - the elements after which the building of a new quantum level (and a new period) begins - on the outer quantum layer contains only 8 electrons and have numbers 2, 10, 18, 36, 54, and 86. Accordingly, the periods cover 2, 8, 8, 18, 18 and 32 elements.

The periodic law has no definite mathematical expression. It is presented in the form of a periodic table. There are several variants of such a table, but all of them in one form or another are presented as structural diagrams of the atomic structure of any element. It becomes possible to establish electronic structure any atom not only on the basis of the known sequence of filling the sublevels or the Klechkovsky rule, but also on the basis of the table itself: the position of an element in the table unambiguously reflects the electronic structure of its atoms. The distribution of elements by periods and by subgroups exactly corresponds to the distribution of electrons of atoms of these elements over the levels and sublevels of the electron shell.

2.3. DI Mendeleev's periodic law.

The law was discovered and formulated by DI Mendeleev: "The properties of simple bodies, as well as the forms and properties of compounds of elements are periodically dependent on the atomic weights of the elements." The law was created on the basis of a deep analysis of the properties of elements and their compounds. Outstanding achievement physicists, mainly the development of the theory of the structure of the atom, made it possible to reveal the physical essence of the periodic law: the periodicity in the change in the properties of chemical elements is due to a periodic change in the nature of filling the outer electron layer with electrons as the number of electrons increases, determined by the nuclear charge. The charge is equal to the ordinal number of the element in the periodic system. The modern formulation of the periodic law: "Properties of elements and the simple and complex substances are periodically dependent on the charge of the atomic nucleus ”. Created by D.I. Mendeleev in 1869-1871. the periodic system is a natural classification of elements, a mathematical reflection of the periodic law.

Mendeleev was not only the first to accurately formulate this law and present its content in the form of a table, which became classical, but also comprehensively substantiated it, showed its enormous scientific significance, as a guiding classification principle and as a powerful instrument for scientific research.

The physical meaning of the periodic law. It was discovered only after it was found out that the charge of the atomic nucleus increases during the transition from one chemical element to the neighboring one (in the periodic system) per unit of elementary charge. Numerically, the charge of the nucleus is equal to the ordinal number (atomic number Z) of the corresponding element in the periodic system, that is, the number of protons in the nucleus, in turn equal to the number of electrons of the corresponding neutral atom. The chemical properties of atoms are determined by the structure of their outer electron shells, which periodically change with an increase in the nuclear charge, and, therefore, the basis of the periodic law is the idea of \u200b\u200ba change in the charge of the nucleus of atoms, not the atomic mass of elements. A clear illustration of the periodic law is the curves of periodic changes in some physical quantities (ionization potentials, atomic radii, atomic volumes) depending on Z. There is no general mathematical expression for the periodic law. The periodic law is of great natural science and philosophical significance. It made it possible to consider all the elements in their interconnection and predict the properties of unknown elements. Thanks to the periodic law, many scientific searches (for example, in the field of studying the structure of matter - in chemistry, physics, geochemistry, cosmochemistry, astrophysics) have become purposeful. The periodic law is a vivid manifestation of the operation of the general laws of dialectics, in particular the law of the transition from quantity to quality.

The physical stage of development of the periodic law can, in turn, be divided into several stages:

1. Establishment of the divisibility of the atom based on the discovery of the electron and radioactivity (1896-1897);

2. Development of models of the structure of the atom (1911-1913);

3. Discovery and development of the isotope system (1913);

4. Discovery of Moseley's law (1913), which makes it possible to experimentally determine the nuclear charge and the number of an element in the periodic table;

5. Development of the theory of the periodic system based on ideas about the structure of the electron shells of atoms (1921-1925);

6. Creation of the quantum theory of the periodic system (1926-1932).


2.4. Predicting the existence of unknown elements.

The most important thing in the discovery of the Periodic Law is the prediction of the existence of not yet discovered chemical elements. Under aluminum Al Mendeleev left a place for its analogue "ekaaluminium", under boron B - for "ekabor", and under silicon Si - for "ecasilicon". So named Mendeleev not yet discovered chemical elements. He even gave them the symbols El, Eb and Es.

Regarding the element "ekasilitsiya" Mendeleev wrote: "It seems to me that the most interesting of the undoubtedly missing metals will be the one that belongs to the IV group of carbon analogs, namely, the III row. This will be the metal immediately following silicon, and therefore we will call it ecasilicon. " Indeed, this not yet discovered element was supposed to become a kind of "lock" connecting two typical non-metals - carbon C and silicon Si - with two typical metals - tin Sn and lead Pb.

Then he predicted the existence of eight more elements, including "dvitellura" - polonium (discovered in 1898), "ekaiod" - astatine (discovered in 1942-1943), "dvimarganese" - technetium (discovered in 1937) , "ecatsia" - France (opened in 1939)

In 1875, the French chemist Paul-Emile Lecoq de Boisbaudran discovered in the mineral wurtzite - zinc sulfide ZnS - predicted by Mendeleev "ekaaluminium" and named it in honor of his homeland gallium Ga (the Latin name for France is "Gaul").

Mendeleev accurately predicted the properties of eka-aluminum: its atomic mass, metal density, the formula of the oxide El 2 O 3, chloride ElCl 3, sulfate El 2 (SO 4) 3. After the discovery of gallium, these formulas began to be written as Ga 2 O 3, GaCl 3 and Ga 2 (SO 4) 3. Mendeleev predicted that it would be a very low-melting metal, and indeed, the melting point of gallium turned out to be equal to 29.8 o C. In terms of low melting point, gallium is second only to mercury Hg and cesium Cs.

The average gallium content in the earth's crust is relatively high, 1.5-10-30% by weight, which is equal to the content of lead and molybdenum. Gallium is a typical trace element. The only gallium mineral, galdite CuGaS2, is very rare. Gallium is stable in air at ordinary temperatures. Above 260 ° C in dry oxygen, slow oxidation is observed (the oxide film protects the metal). In the sulfur and hydrochloric acids gallium dissolves slowly, in hydrofluoric acid it dissolves quickly, in nitric acid in the cold Gallium is stable. Gallium slowly dissolves in hot alkali solutions. Chlorine and bromine react with gallium in the cold, iodine - when heated. Molten Gallium at temperatures above 300 ° C interacts with all structural metals and alloys A distinctive feature of Gallium is a large interval of liquid state (2200 ° C) and low vapor pressure at temperatures up to 1100-1200 ° C. Geochemistry of Gallium is closely related to the geochemistry of aluminum, which is due to the similarity of their physical and chemical properties. The bulk of gallium in the lithosphere is contained in aluminum minerals. The gallium content in bauxite and nepheline ranges from 0.002 to 0.01%. Increased concentrations of gallium are also observed in sphalerites (0.01-0.02%), in coal (together with germanium), and also in some iron ores. Gallium does not yet have wide industrial application. The potentially possible scale of the associated production of gallium in aluminum production still significantly exceeds the demand for the metal.

The most promising application of gallium is in the form of chemical compounds such as GaAs, GaP, GaSb with semiconducting properties. They can be used in high-temperature rectifiers and transistors, solar batteries and other devices, where the photoelectric effect can be used in the blocking layer, as well as in infrared receivers. Gallium can be used to make highly reflective optical mirrors. An alloy of aluminum with Gallium has been proposed instead of mercury as the cathode of ultraviolet radiation lamps used in medicine. Liquid gallium and its alloys have been proposed to be used for the manufacture of high-temperature thermometers (600-1300 ° C) and manometers. Of interest is the use of Gallium and its alloys as a liquid coolant in nuclear power reactors (this is hindered by the active interaction of Gallium at operating temperatures with structural materials; the Ga-Zn-Sn eutectic alloy has a lower corrosive effect than pure Gallium).

In 1879, Swedish chemist Lars Nilsson discovered scandium, predicted by Mendeleev as ekabor Eb. Nilsson wrote: "There is no doubt that an ekabor has been discovered in scandium ... This clearly confirms the considerations of the Russian chemist, which not only made it possible to predict the existence of scandium and gallium, but also to foresee their most important properties in advance." Scandium was named after Nielson's homeland of Scandinavia, and he discovered it in the complex mineral gadolinite, which has the composition Be 2 (Y, Sc) 2 FeO 2 (SiO 4) 2. The average content of Scandium in the earth's crust (clarke) is 2.2-10-3% by weight. In rocks, the content of Scandium is different: in ultrabasic 5-10-4, in basic 2.4-10-3, in middle 2.5-10-4, in granites and syenites 3.10-4; in sedimentary rocks (1-1.3) .10-4. Scandium is concentrated in the earth's crust as a result of magmatic, hydrothermal and hypergene (surface) processes. There are two known Scandium minerals of their own - tortveitite and sterrettite; they are extremely rare. Scandium is a soft metal; in its pure state it can be easily processed - forging, rolling, stamping. Scandium is very limited in scope. Scandium oxide is used to make ferrites for memory elements in high-speed computers. The radioactive 46Sc is used in neutron activation analysis and in medicine. Scandium alloys, which have a low density and high melting point, are promising as structural materials in rocket and aircraft construction, and a number of Scandium compounds can be used in the manufacture of phosphors, oxide cathodes, in glass and ceramic industries, in the chemical industry (as catalysts) and others. areas. In 1886, a professor at the Mining Academy in Freiburg, German chemist Clemens Winkler, while analyzing a rare mineral argyrodite of the composition Ag 8 GeS 6, discovered another element predicted by Mendeleev. Winkler named the element Ge, which he discovered, after his homeland, but for some reason this provoked strong objections from some chemists. They began to accuse Winkler of nationalism, of appropriating the discovery made by Mendeleev, who had already given the element the name "ekasiliciy" and the symbol Es. Discouraged, Winkler turned to Dmitry Ivanovich himself for advice. He explained that it was the discoverer of the new element that should give it a name. The total content of Germanium in the earth's crust is 7.10-4% by mass, that is, more than, for example, antimony, silver, bismuth. However, the native minerals Germanium are extremely rare. Almost all of them are sulfosalts: germanite Cu2 (Cu, Fe, Ge, Zn) 2 (S, As) 4, argyrodite Ag8GeS6, confildite Ag8 (Sn, Ce) S6, etc. The bulk of Germanium is scattered in the earth's crust in large numbers rocks and minerals: in sulfide ores of non-ferrous metals, in iron ores, in some oxide minerals (chromite, magnetite, rutile, etc.), in granites, diabases and basalts. In addition, Germanium is present in almost all silicates, in some deposits of coal and oil. Germanium is one of the most valuable materials in modern semiconductor technology. It is used to make diodes, triodes, crystal detectors and power rectifiers. Monocrystalline Germanium is also used in dosimetry devices and devices that measure the strength of constant and alternating magnetic fields. An important field of application of germanium is infrared technology, in particular the production of infrared detectors operating in the 8-14 micron range. Perspective for practical use many alloys, which include Germanium, glasses based on GeO2, and other compounds of Germanium.

Mendeleev could not predict the existence of a group of noble gases, and at first there was no place for them in the Periodic Table.

The discovery of argon Ar by the English scientists W. Ramsay and J. Rayleigh in 1894 immediately caused heated discussions and doubts about the Periodic Law and the Periodic Table of Elements. Mendeleev initially considered argon to be an allotropic modification of nitrogen, and only in 1900, under the pressure of immutable facts, agreed with the presence in the Periodic Table of the "zero" group of chemical elements, which was occupied by other noble gases that were discovered after argon. This group is now known under the number VIIIA.

In 1905, Mendeleev wrote: "Apparently, the future does not threaten the periodic law with destruction, but only promises superstructures and development, although they wanted to wipe me out as a Russian, especially the Germans."

The discovery of the Periodic Law hastened the development of chemistry and the discovery of new chemical elements.

Lyceum exam, at which old man Derzhavin blessed young Pushkin. Academician Yu.F. Fritzsche, a well-known specialist in organic chemistry, had a chance to play the role of meter. D. I. Mendeleev graduated from the Main Pedagogical Institute in 1855. The candidate thesis "Isomorphism in connection with other relations of the crystalline form to the composition" became his first major scientific ...

Mainly on the issue of capillarity and surface tension of liquids, and spent his leisure hours in the circle of young Russian scientists: S.P. Botkin, I.M. Sechenov, I.A. Vyshnegradsky, A.P. Borodin and others. In 1861 Mendeleev returned to St. Petersburg, where he resumed lecturing on organic chemistry at the university and published a textbook, remarkable for that time: " Organic chemistry", in...

DI Mendeleev's periodic law, its modern formulation. What is its difference from the one given by D.I. Mendeleev? Explain what caused such a change in the wording of the law? What is the physical meaning of the Periodic Law? Explain the reason for the periodic change in the properties of chemical elements. How do you understand the phenomenon of periodicity?

The periodic law was formulated by DI Mendeleev in the following form (1871): "the properties of simple bodies, as well as the forms and properties of compounds of elements, and therefore the properties of simple and complex bodies formed by them, are periodically dependent on their atomic weight."

At present, DI Mendeleev's Periodic Law has the following formulation: "the properties of chemical elements, as well as the forms and properties of the simple substances and compounds formed by them, are periodically dependent on the magnitude of the charges of the nuclei of their atoms."

The peculiarity of the Periodic Law among other fundamental laws is that it has no expression in the form of a mathematical equation. The graphic (tabular) expression of the law is the Periodic Table of Elements developed by Mendeleev.

The periodic law is universal for the Universe: as the famous Russian chemist ND Zelinsky figuratively noted, the periodic law was "the discovery of the mutual connection of all atoms in the universe."

IN current state The periodic table of elements consists of 10 horizontal rows (periods) and 8 vertical columns (groups). The first three rows form three small periods. Subsequent periods include two rows. In addition, starting from the sixth, the periods include additional series of lanthanides (sixth period) and actinides (seventh period).

Over the period, there is a weakening of metallic properties and an increase in non-metallic properties. The end element of the period is a noble gas. Each subsequent period begins with an alkali metal, i.e., as the atomic mass of elements grows, the change in chemical properties has a periodic character.

With the development of atomic physics and quantum chemistry, the Periodic Law received a rigorous theoretical foundation. Thanks to the classical works of J. Rydberg (1897), A. Van den Bruck (1911), G. Moseley (1913), the physical meaning of the ordinal (atomic) number of an element was revealed. Later, a quantum-mechanical model was created for the periodic change in the electronic structure of atoms of chemical elements as the charges of their nuclei increased (N. Bohr, W. Pauli, E. Schrödinger, W. Heisenberg, and others).

Periodic properties of chemical elements

In principle, the properties of a chemical element unite all, without exception, its characteristics in the state of free atoms or ions, hydrated or solvated, in the state of a simple substance, as well as the forms and properties of the numerous compounds formed by it. But usually the properties of a chemical element mean, firstly, the properties of its free atoms and, secondly, the properties of a simple substance. Most of these properties show a clear periodic dependence on the atomic numbers of chemical elements. Among these properties, the most important ones that are of particular importance in explaining or predicting the chemical behavior of elements and the compounds they form are:

Ionization energy of atoms;

The energy of the affinity of atoms for an electron;

Electronegativity;

Atomic (and ionic) radii;

Atomization energy of simple substances

Oxidation states;

Oxidizing Potentials of Simple Substances.

The physical meaning of the periodic law is that the periodic change in the properties of elements is in full accordance with the periodically renewed at ever higher energy levels similar electronic structures of atoms. With their regular change, physical and chemical properties change naturally.

The physical meaning of the periodic law became clear after the creation of the theory of the structure of the atom.

So, the physical meaning of the periodic law is that the periodic change in the properties of the elements is in full accordance with the periodically renewed at higher and higher energy levels similar electronic structures of atoms. With their regular change, the physical and chemical properties of the elements change naturally.

What is the physical meaning of the periodic law.

These conclusions reveal the physical meaning of the periodic law of D. I. Mendeleev, which remained unclear for half a century after the discovery of this law.

It follows from this that the physical meaning of the periodic law of D.I.

The theory of the structure of atoms has shown that the physical meaning of the periodic law is that with a successive increase in the charges of nuclei, similar valence electronic structures of atoms are periodically repeated.

From all that has been said, it is clear that the theory of atomic structure revealed the physical meaning of the periodic law of D.I.Mendeleev and even more clearly revealed its significance as the basis for the further development of chemistry, physics and a number of other sciences.

Replacing the atomic mass with the nuclear charge was the first step in revealing the physical meaning of the periodic law. Further, it was important to establish the reasons for the occurrence of periodicity, the nature periodic function dependence of properties on the charge of the nucleus, explain the values \u200b\u200bof the periods, the number of rare-earth elements, etc.

For analog elements, the same number of electrons is observed on the shells of the same name at different meanings the principal quantum number. Therefore, the physical meaning of the Periodic Law lies in the periodic change in the properties of elements as a result of periodically renewing similar electron shells of atoms with a sequential increase in the values \u200b\u200bof the principal quantum number.

For elements - analogs, the same number of electrons is observed in the same orbitals at different values \u200b\u200bof the principal quantum number. Therefore, the physical meaning of the Periodic Law lies in the periodic change in the properties of elements as a result of periodically renewing similar electron shells of atoms with a sequential increase in the values \u200b\u200bof the principal quantum number.

Thus, with a sequential increase in the charges of atomic nuclei, the configuration of the electron shells is periodically repeated and, as a consequence, the chemical properties of the elements are periodically repeated. This is the physical meaning of the periodic law.

DI Mendeleev's periodic law is the basis of modern chemistry. The study of the structure of atoms reveals the physical meaning of the periodic law and explains the regularities of changes in the properties of elements in periods and in groups of the periodic system. Knowledge of the structure of atoms is necessary to understand the reasons for the formation chemical bond... The nature of the chemical bond in molecules determines the properties of substances. Therefore, this section is one of the most important sections of general chemistry.

natural science periodic ecosystem

: as the famous Russian chemist ND Zelinsky figuratively noted, the Periodic Law was "the discovery of the mutual connection of all atoms in the universe."

History

The search for the basis for the natural classification and systematization of chemical elements began long before the discovery of the Periodic Law. The difficulties encountered by the naturalists who were the first to work in this area were caused by the lack of experimental data: at the beginning of the 19th century, the number of known chemical elements was small, and the accepted values \u200b\u200bof the atomic masses of many elements are incorrect.

Doebereiner's triads and the first systems of elements

In the early 60s of the XIX century, several works appeared at once, which immediately preceded the Periodic Law.

Spiral de Chancourtois

Octaves of Newlands

Newlands Table (1866)

Soon after de Chancourtois spiral, the English scientist John Newlands made an attempt to compare the chemical properties of elements with their atomic masses. Arranging the elements in ascending order of their atomic masses, Newlands noticed that similarities in properties appear between every eighth element. The found pattern Newlands called the law of octaves by analogy with seven intervals of a musical scale. In his table, he arranged the chemical elements in vertical groups of seven elements in each and at the same time found that (with a slight change in the order of some elements) similar in chemical properties the elements fall on the same horizontal line.

John Newlands, undoubtedly, was the first to give a number of elements arranged in the order of increasing atomic masses, assigned the corresponding serial number to the chemical elements and noticed a systematic relationship between this order and physical and chemical properties elements. He wrote that in such a sequence the properties of elements are repeated, equivalent weights (masses) of which differ by 7 units, or by a value that is a multiple of 7, that is, as if the eighth in order element repeats the properties of the first, as in music the eighth note repeats first. Newlands endeavored to make this dependence, which is indeed the case for the light elements, universal. In his table, similar elements were located in horizontal rows, but elements with completely different properties were often in the same row. In addition, Newlands had to accommodate two elements in some cells; finally, the table did not contain empty spaces; as a result, the law of octaves was received with extreme skepticism.

Odling and Meier tables

Manifestations of the periodic law in relation to the electron affinity energy

The periodicity of the values \u200b\u200bof the energies of the affinity of atoms for an electron is naturally explained by the same factors that have already been noted in the discussion of ionization potentials (see the definition of the energy of affinity for an electron).

The greatest affinity for the electron is possessed by p- elements of group VII. The smallest electron affinity is for atoms with the s² (,,) and s²p 6 (,) configurations or with half-filled p-orbitals (,,):

Manifestations of the periodic law of electronegativity

Strictly speaking, an element cannot be ascribed permanent electronegativity. The electronegativity of an atom depends on many factors, in particular, on the valence state of the atom, the formal oxidation state, the coordination number, the nature of the ligands that make up the environment of the atom in the molecular system, and some others. Recently, the so-called orbital electronegativity has been increasingly used to characterize electronegativity, which depends on the type of atomic orbital involved in the formation of a bond and on its electronic population, i.e. on whether the atomic orbital is occupied by a lone electron pair, is populated once by an unpaired electron, or is vacant. But, despite the well-known difficulties in the interpretation and definition of electronegativity, it always remains necessary for a qualitative description and prediction of the nature of bonds in a molecular system, including the binding energy, distribution of electronic charge and degree of ionicity, force constant, etc.

The periodicity of atomic electronegativity is an important component of the periodic law and can be easily explained proceeding from the immutable, although not entirely unambiguous, dependence of the electronegativity values \u200b\u200bon the corresponding values \u200b\u200bof ionization energies and electron affinity.

In periods, there is a general tendency for the growth of electronegativity, and in subgroups, its decline. The smallest electronegativity is for s-elements of group I, the highest is for p-elements of group VII.

Manifestations of the periodic law in relation to atomic and ionic radii

Figure: 4 Dependence of the orbital radii of atoms on the ordinal number of the element.

The periodic nature of changes in the size of atoms and ions has long been known. The difficulty here lies in the fact that, due to the wave nature of the electronic motion, atoms do not have strictly defined sizes. Since a direct determination of the absolute sizes (radii) of isolated atoms is impossible, in this case their empirical values \u200b\u200bare often used. They are obtained from the measured internuclear distances in crystals and free molecules, breaking each internuclear distance into two parts and equating one of them to the radius of the first (of two connected by the corresponding chemical bond) atom, and the other to the radius of the second atom. In this division, various factors are taken into account, including the nature of the chemical bond, the oxidation state of the two bound atoms, the nature of the coordination of each of them, etc. In this way, the so-called metallic, covalent, ionic and van der Waals radii are obtained. Van der Waals radii should be considered as the radii of unbound atoms; they are found by the internuclear distances in solid or liquid substances, where atoms are in close proximity to each other (for example, atoms in solid argon or atoms from two neighboring N 2 molecules in solid nitrogen), but are not linked to each other by any chemical bond ...

But, obviously, the best description of the effective dimensions of an isolated atom is the theoretically calculated position (distance from the nucleus) of the main maximum of the charge density of its outer electrons. This is the so-called orbital radius of the atom. The periodicity in the change in the values \u200b\u200bof the orbital atomic radii, depending on the ordinal number of the element, manifests itself quite clearly (see Fig. 4), and the main points here are the presence of very pronounced maxima corresponding to atoms of alkali metals, and the same minima corresponding to noble gases ... A decrease in the values \u200b\u200bof orbital atomic radii during the transition from an alkali metal to the corresponding (nearest) noble gas is, with the exception of the - series, non-monotonic, especially when families of transition elements (metals) and lanthanides or actinides appear between the alkali metal and the noble gas. In large periods in families d- and f-elements, a less sharp decrease in radii is observed, since the filling of the orbitals with electrons occurs in the pre-outer layer. In subgroups of elements, the radii of atoms and ions of the same type generally increase.

Manifestations of the periodic law in relation to atomization energy

It should be emphasized that the oxidation state of an element, being a formal characteristic, does not give an idea of \u200b\u200beither the effective charges of the atoms of this element in the compound, or the valence of atoms, although the oxidation state is often called the formal valence. Many elements are capable of exhibiting not one but several different oxidation states. For example, for chlorine, all oxidation states are known from −1 to +7, although even ones are very unstable, and for manganese, from +2 to +7. The highest values \u200b\u200bof the oxidation state change periodically depending on the ordinal number of the element, but this periodicity is complex. In the simplest case, in the series of elements from an alkali metal to a noble gas, the highest oxidation state increases from +1 (F) to +8 (O 4). In other cases, the highest oxidation state of the noble gas is lower (+4 F 4) than for the preceding halogen (+7 O 4 -). Therefore, on the curve of the periodic dependence of the highest oxidation state on the ordinal number of the element, the maxima fall either on the noble gas or on the halogen preceding it (the minima are always on the alkali metal). The exception is the series -, in which high oxidation states are generally unknown neither for halogen (), nor for noble gas (), and the middle term of the series, nitrogen, has the highest value of the highest oxidation state; therefore, in the series - the change in the highest oxidation state turns out to be passing through a maximum. In the general case, the increase in the highest oxidation state in the series of elements from an alkali metal to a halogen or to a noble gas is by no means monotonic, mainly due to the manifestation of high oxidation states transition metals... For example, an increase in the highest oxidation state in the series - from +1 to +8 is "complicated" by the fact that for molybdenum, technetium and ruthenium such high oxidation states as +6 (О 3), +7 (2 О 7), + 8 (O 4).

Manifestations of the periodic law in relation to oxidative potential

One of the very important characteristics of a simple substance is its oxidative potential, which reflects the fundamental ability of a simple substance to interact with aqueous solutions, as well as the redox properties shown by it. The change in the oxidative potentials of simple substances depending on the ordinal number of the element is also periodic. But it should be borne in mind that the oxidative potential of a simple substance is influenced by various factors, which sometimes need to be considered individually. Therefore, the periodicity in the change in oxidation potentials should be interpreted very carefully.

/ Na + (aq) / Mg 2+ (aq) / Al 3+ (aq)
2.71V 2.37V 1.66V
/ K + (aq) / Ca 2+ (aq) / Sc 3+ (aq)
2.93V 2.87V 2.08V

You can find some specific sequences in the change in the oxidative potentials of simple substances. In particular, in the series of metals, when passing from alkaline to the elements following it, the oxidation potentials decrease (+ (aq), etc. - hydrated cation):

This is easily explained by an increase in the ionization energy of atoms with an increase in the number of removed valence electrons. Therefore, the curve of the dependence of the oxidation potentials of simple substances on the ordinal number of the element has maxima corresponding to alkali metals. But this is not the only reason for the change in the oxidative potentials of simple substances.

Internal and secondary periodicity

s- and r-elements

Above, general trends in the nature of changes in the values \u200b\u200bof the ionization energy of atoms, the energy of the affinity of atoms to the electron, electronegativity, atomic and ionic radii, the energy of atomization of simple substances, the oxidation state, oxidation potentials of simple substances from the atomic number of an element are considered. With a deeper study of these trends, one can find that the patterns in the change in the properties of elements in periods and groups are much more complicated. In the nature of the change in the properties of elements by period, internal periodicity is manifested, and in the group - secondary periodicity (discovered by E.V. Biron in 1915).

Thus, when passing from an s-element of group I to r-element of group VIII on the curve of the ionization energy of atoms and the curve of change of their radii have internal maxima and minima (see Fig. 1, 2, 4).

This indicates the internally periodic nature of changes in these properties over the period. The above regularities can be explained using the concept of nuclear screening.

The shielding effect of the nucleus is caused by the electrons of the inner layers, which, blocking the nucleus, weaken the attraction of the outer electron to it. So, when going from beryllium 4 to boron 5, despite an increase in the nuclear charge, the ionization energy of atoms decreases:

Figure: 5 Diagram of the structure of the last levels of beryllium, 9.32 eV (left) and boron, 8.29 eV (right)

This is because the attraction to the core 2p-electron of boron atom is weakened due to the shielding action 2s-electrons.

It is clear that the screening of the nucleus increases with an increase in the number of internal electron layers. Therefore, in the subgroups s- and r-elements, there is a tendency to a decrease in the ionization energy of atoms (see Fig. 1).

The decrease in the ionization energy from nitrogen 7 N to oxygen 8 O (see Fig. 1) is explained by the mutual repulsion of two electrons of the same orbital:

Figure: 6 Scheme of the structure of the last levels of nitrogen, 14.53 eV (left) and oxygen, 13.62 eV (right)

The effect of screening and mutual repulsion of electrons of one orbital also explains the internally periodic nature of the change over the period of atomic radii (see Fig. 4).

Figure: 7 Secondary periodic dependence of the radii of the atoms of the outer p-orbitals on the atomic number

Figure: 8 Secondary periodic dependence of the first ionization energy of atoms on the atomic number

Figure: 9 Radial distribution of electron density in the sodium atom

In the nature of the change in properties s- and r-elements in subgroups, a secondary periodicity is clearly observed (Fig. 7). To explain it, the concept of electron penetration to the nucleus is used. As shown in Figure 9, an electron of any orbital for a certain time is in a region close to the nucleus. In other words, the outer electrons penetrate to the nucleus through the layers of inner electrons. As can be seen from Figure 9, external 3 s-electron of a sodium atom has a very significant probability of being near the nucleus in the region of internal TO- and L-electronic layers.

The concentration of electron density (the degree of penetration of electrons) at the same principal quantum number is highest for s-electron, less - for r-electron, even less - for d-electron, etc. For example, for n \u003d 3, the degree of penetration decreases in the sequence 3 s>3p>3d (see fig. 10).

Figure: 10 Radial distribution of the probability of finding an electron (electron density) at a distance r from the core

It is clear that the penetration effect increases the strength of the bond between the outer electrons and the nucleus. Due to deeper penetration s-electrons screen the nucleus to a greater extent than r-electrons, and the latter are stronger than d-electrons, etc.

Using the concept of electron penetration to the nucleus, let us consider the nature of the change in the radius of the atoms of the elements in the carbon subgroup. In the series - - - - there is a general tendency to increase the radius of the atom (see Fig. 4, 7). However, this increase is non-monotonic. On going from Si to Ge, the external r- electrons penetrate the screen out of ten 3 d-electrons and thereby strengthen the bond with the nucleus and compress the electron shell of the atom. Downsize 6 p-orbitals of Pb compared to 5 r-orbital Sn is due to the penetration of 6 p-electrons under the double screen ten 5 d-electrons and fourteen 4 f-electrons. This also explains the nonmonotonicity in the change in the ionization energy of atoms in the C-Pb series and its greater value for Pb in comparison with the Sn atom (see Fig. 1).

d-Elements

In the outer layer of atoms d-elements (except for) there are 1-2 electrons ( ns-state). The rest of the valence electrons are located in (n-1) d-state, that is, in the pre-outer layer.

A similar structure of the electronic shells of atoms determines some general properties d-elements. Thus, their atoms are characterized by relatively low values \u200b\u200bof the first ionization energy. As can be seen in Fig. 1, the nature of the change in the ionization energy of atoms over the period in the series d- the elements are smoother than in a row s- and p-elements. When going from d-element of III group to d-element of the II group, the values \u200b\u200bof the ionization energy change non-monotonically. Thus, in the segment of the curve (Fig. 1), two areas are visible, corresponding to the ionization energy of atoms, in which 3 d-orbitals one and two electrons each. Filling 3 d-orbitals, one electron each ends at (3d 5 4s 2), which is marked by a slight increase in the relative stability of the 4s 2 -configuration due to the penetration of 4s 2 -electrons under the shield of the 3d 5 -configuration. The highest value of the ionization energy has (3d 10 4s 2), which is in accordance with the complete completion of 3 d-sublayer and stabilization of the electron pair due to penetration under the screen 3 d 10 configurations.

In subgroups d-elements, the values \u200b\u200bof the ionization energy of atoms generally increase. This can be explained by the effect of electron penetration to the nucleus. So if u d- elements of the 4th period external 4 s-electrons penetrate the screen 3 d-electrons, then the elements of the 6th period have external 6 s-electrons already penetrate under the double screen 5 d- and 4 f-electrons. For example:

22 Ti ... 3d 2 4s 2 I \u003d 6.82 eV
40 Zr… 3d 10 4s 2 4p 6 4d 2 5s 2 I \u003d 6.84 eV
72 Hf… 4d 10 4f 14 5s 2 5p 6 5d 2 6s 2 I \u003d 7.5 eV

Therefore, d-elements of the 6th period external b s-electrons are bound to the nucleus more firmly and, therefore, the ionization energy of atoms is greater than that of d-elements of the 4th period.

Dimensions of atoms d-elements are intermediate between the sizes of atoms s- and p- elements of this period. The change in the radii of their atoms over the period is smoother than for s- and p-elements.

In subgroups d-elements, the radii of the atoms generally increase. It is important to note the following feature: an increase in atomic and ionic radii in subgroups d-elements basically corresponds to the transition from the 4th element to the 5th period element. The corresponding radii of atoms d-elements of the 5th and 6th periods of this subgroup are approximately the same. This is explained by the fact that an increase in radii due to an increase in the number of electron layers during the transition from the 5th to the 6th period is compensated for by f-compression caused by filling with electrons 4 f-sublayer at f-elements of the 6th period. In this case f-compression is called lanthanoid... With similar electronic configurations of the outer layers and approximately the same sizes of atoms and ions for d-elements of the 5th and 6th periods of this subgroup are characterized by a special similarity of properties.

The elements of the scandium subgroup do not obey these patterns. This subgroup is characterized by patterns characteristic of neighboring subgroups s-elements.

The periodic law is the basis of chemical taxonomy

see also

Notes

Literature

  1. Akhmetov N.S. Topical issues of the course inorganic chemistry... - M .: Education, 1991 .-- 224 p. - ISBN 5-09-002630-0
  2. Korolkov D.V. Fundamentals of Inorganic Chemistry. - M .: Education, 1982 .-- 271 p.
  3. Mendeleev D.I. Fundamentals of Chemistry, vol. 2. M .: Goskhimizdat, 1947.389 p.
  4. Mendeleev D.I. // Encyclopedic Dictionary of Brockhaus and Efron: In 86 volumes (82 volumes and 4 additional). - SPb. , 1890-1907.