The prerequisite for the discovery of the Periodic Law was the decisions of the international congress of chemists in the city of Karlsruhe in 1860, when the atomic-molecular science was finally established and the first unified definitions of the concepts of molecule and atom, as well as atomic weight, which we now call relative atomic mass, were undertaken.
D.I. Mendeleev in his discovery relied on clearly formulated starting points:
General unchanging property of all atoms chemical elements- their atomic mass;
The properties of elements depend on their atomic masses;
The form of this dependence is periodic.
The prerequisites discussed above can be called objective, that is, independent of the personality of the scientist, since they were determined by the historical development of chemistry as a science.
III Periodic Law and Periodic Table of Chemical Elements.
Mendeleev's discovery of the Periodic Law.
The first version of the Periodic Table of Elements was published by D. I. Mendeleev in 1869 - long before the structure of the atom was studied. At this time, Mendeleev taught chemistry at St. Petersburg University. Preparing for lectures, collecting material for his textbook “Fundamentals of Chemistry,” D. I. Mendeleev thought about how to systematize the material in such a way that information about chemical properties ah elements did not look like a set of disparate facts.
D. I. Mendeleev’s guide in this work was the atomic masses (atomic weights) of elements. After the World Congress of Chemists in 1860, in which D.I. Mendeleev also participated, the problem of correctly determining atomic weights was constantly in the focus of attention of many leading chemists in the world, including D.I. Mendeleev.By arranging the elements in increasing order of their atomic weights, D. I. Mendeleev discovered a fundamental law of nature, which is now known as the Periodic Law:
The properties of elements change periodically according to their atomic weight.
The above formulation does not at all contradict the modern one, in which the concept of “atomic weight” is replaced by the concept of “nuclear charge”. The nucleus consists of protons and neutrons. The number of protons and neutrons in the nuclei of most elements is approximately the same, so atomic weight increases in approximately the same way as the number of protons in the nucleus (nuclear charge Z) increases.
The fundamental novelty of the Periodic Law was as follows:
1. A connection was established between elements that were dissimilar in their properties. This connection lies in the fact that the properties of elements change smoothly and approximately equally as their atomic weight increases, and then these changes REPEAT PERIODICALLY.
2. In those cases when it seemed that some link was missing in the sequence of changes in the properties of elements, GAPS were provided in the Periodic Table that had to be filled with elements that had not yet been discovered.
In all previous attempts to determine the relationship between elements, other researchers sought to create a complete picture in which there was no room for elements that had not yet been discovered. On the contrary, D.I. Mendeleev considered the most important part of his Periodic Table to be those cells that were still empty. This made it possible to predict the existence of still unknown elements.
It is admirable that D. I. Mendeleev made his discovery at a time when the atomic weights of many elements were determined very approximately, and only 63 elements themselves were known - that is, a little more than half of those known to us today.
Deep knowledge of the chemical properties of various elements allowed Mendeleev not only to point out elements that had not yet been discovered, but also to accurately predict their properties! D.I. Mendeleev accurately predicted the properties of the element he called “eka-silicon”. 16 years later, this element was indeed discovered by the German chemist Winkler and named germanium.
Comparison of the properties predicted by D.I. Mendeleev for the yet undiscovered element “eka-silicon” with the properties of the element germanium (Ge). In the modern Periodic Table, germanium occupies the place of "eka-silicon".
Property
Predicted by D.I. Mendeleev for "eka-silicon" in 1870
Defined for germanium Ge, discovered in 1886
Color, appearance
brown
light brown
Atomic weight
72,59
Density (g/cm3)
5,5
5,35
Oxide formula
XO2
GeO2
Chloride formula
XCl4
GeCl4
Chloride Density (g/cm3)
1,9
1,84
In the same way, the properties of “eka-aluminum” (the element gallium Ga, discovered in 1875) and “eka-boron” (the element scandium Sc, discovered in 1879) were brilliantly confirmed by D.I. Mendeleev.
After this, it became clear to scientists around the world that D. I. Mendeleev’s Periodic Table does not simply systematize the elements, but is a graphic expression of the fundamental law of nature - the Periodic Law.
Structure of the Periodic Table.
Based on the Periodic Law of D.I. Mendeleev created the Periodic Table of Chemical Elements, which consisted of 7 periods and 8 groups (short-period version of the table). Currently, the long-period version of the Periodic System is more often used (7 periods, 8 groups, the elements lanthanides and actinides are shown separately).
Periods are horizontal rows of the table; they are divided into small and large. In small periods there are 2 elements (1st period) or 8 elements (2nd, 3rd periods), in large periods - 18 elements (4th, 5th periods) or 32 elements (6th, 5th periods) 7th period). Each period begins with a typical metal and ends with a nonmetal (halogen) and a noble gas.
Groups are vertical sequences of elements, they are numbered with Roman numerals from I to VIII and Russian letters A and B. The short-period version of the Periodic System included subgroups of elements (main and secondary).
A subgroup is a set of elements that are unconditional chemical analogues; often elements of a subgroup have the highest oxidation state corresponding to the group number.
In A-groups, the chemical properties of elements can vary over a wide range from non-metallic to metallic (for example, in the main subgroup of group V, nitrogen is a non-metal, and bismuth is a metal).
In the Periodic Table, typical metals are located in group IA (Li-Fr), IIA (Mg-Ra) and IIIA (In, Tl). Non-metals are located in groups VIIA (F-Al), VIA (O-Te), VA (N-As), IVA (C, Si) and IIIA (B). Some elements of A-groups (beryllium Be, aluminum Al, germanium Ge, antimony Sb, polonium Po and others), as well as many elements of B-groups exhibit both metallic and non-metallic properties (the phenomenon of amphotericity).
For some groups, group names are used: IA (Li-Fr) - alkali metals, IIA (Ca-Ra) - alkaline earth metals, VIA (O-Po) - chalcogens, VIIA (F-At) - halogens, VIIIA (He-Rn ) - noble gases. The form of the Periodic Table proposed by D.I. Mendeleev, was called short-period or classical. Currently, another form of the Periodic Table is more widely used - the long-period one.
Periodic law D.I. Mendeleev and the Periodic Table of Chemical Elements became the basis of modern chemistry. Relative atomic masses are given according to the 1983 International Table. For elements 104-108, the mass numbers of the longest-lived isotopes are given in square brackets. The names and symbols of elements given in parentheses are not generally accepted.
IV Periodic law and the structure of the atom.
Basic information on the structure of atoms.
IN late XIX- at the beginning of the 20th century, physicists proved that the atom is a complex particle and consists of simpler (elementary) particles. Were discovered:
cathode rays (English physicist J. J. Thomson, 1897), the particles of which are called electrons e− (carry a single negative charge);
natural radioactivity of elements (French scientists - radiochemists A. Becquerel and M. Sklodowska-Curie, physicist Pierre Curie, 1896) and the existence of α-particles (helium nuclei 4He2+);
the presence of a positively charged nucleus at the center of the atom (English physicist and radiochemist E. Rutherford, 1911);
the artificial transformation of one element into another, for example nitrogen into oxygen (E. Rutherford, 1919). From the nucleus of an atom of one element (nitrogen - in Rutherford’s experiment), upon collision with an α-particle, the nucleus of an atom of another element (oxygen) and a new particle carrying a unit positive charge and called a proton (p+, 1H nucleus) were formed.
the presence in the nucleus of an atom of electrically neutral particles - neutrons n0 (English physicist J. Chadwick, 1932).
As a result of the research, it was found that the atom of each element (except 1H) contains protons, neutrons and electrons, with protons and neutrons concentrated in the nucleus of the atom, and electrons on its periphery (in the electron shell).
The number of protons in the nucleus is equal to the number of electrons in the shell of the atom and corresponds to the serial number of this element in the Periodic Table.
The electron shell of an atom is a complex system. It is divided into subshells with different energies (energy levels); the levels, in turn, are divided into sublevels, and the sublevels include atomic orbitals, which can differ in shape and size (denoted by the letters s, p, d, f, etc.).
So, the main characteristic of an atom is not the atomic mass, but the magnitude of the positive charge of the nucleus. This is a more general and accurate characteristic of an atom, and therefore an element. All properties of the element and its position in the periodic table depend on the magnitude of the positive charge of the atomic nucleus. Thus, the atomic number of a chemical element numerically coincides with the charge of the nucleus of its atom. The periodic table of elements is a graphical representation periodic law and reflects the structure of atoms of elements.
The theory of atomic structure explains the periodic changes in the properties of elements. An increase in the positive charge of atomic nuclei from 1 to 110 leads to a periodic repetition of the structural elements of the external energy level in atoms. And since the properties of elements mainly depend on the number of electrons on the external level, they also repeat periodically. This is the physical meaning of the periodic law.
Each period in the periodic system begins with elements whose atoms at the outer level have one s-electron (incomplete outer levels) and therefore exhibit similar properties - they easily give up valence electrons, which determines their metallic character. These are alkali metals - Li, Na, K, Rb, Cs.
The period ends with elements whose atoms at the outer level contain 2 (s2) electrons (in the first period) or 8 (s2p6) electrons (in all subsequent periods), that is, they have a completed external level. These are noble gases He, Ne, Ar, Kr, Xe, which have inert properties.
The periodic table of elements had a great influence on the subsequent development of chemistry.
Dmitry Ivanovich Mendeleev (1834-1907)
Not only was it the first natural classification of chemical elements, showing that they form a harmonious system and are in close connection with each other, but it also became a powerful tool for further research.
At the time when Mendeleev compiled his table based on the periodic law he discovered, many elements were still unknown. Thus, the fourth period element scandium was unknown. In terms of atomic mass, titanium came after calcium, but titanium could not be placed immediately after calcium, since it would fall into the third group, while titanium forms a higher oxide, and according to other properties it should be classified in the fourth group. Therefore, Mendeleev skipped one cell, that is, he left free space between calcium and titanium. On the same basis, in the fourth period, two free cells were left between zinc and arsenic, now occupied by the elements gallium and germanium. There are still empty seats in other rows. Mendeleev was not only convinced that there must be as yet unknown elements that would fill these spaces, but he also predicted the properties of such elements in advance based on their position among other elements of the periodic table. He gave the name ekabor to one of them, which in the future was to take a place between calcium and titanium (since its properties were supposed to resemble boron); the other two, for which there were spaces left in the table between zinc and arsenic, were named eka-aluminum and eca-silicon.
Over the next 15 years, Mendeleev's predictions were brilliantly confirmed: all three expected elements were discovered. First, the French chemist Lecoq de Boisbaudran discovered gallium, which has all the properties of eka-aluminium; then, in Sweden, L. F. Nilsson discovered scandium, which had the properties of ekaboron, and finally, a few years later in Germany, K. A. Winkler discovered an element he called germanium, which turned out to be identical to ekasilicon.
To judge the amazing accuracy of Mendeleev’s foresight, let us compare the properties of eca-silicon predicted by him in 1871 with the properties of germanium discovered in 1886:
The discovery of gallium, scandium and germanium was the greatest triumph of the periodic law.
The periodic system was also of great importance in establishing the valence and atomic masses of some elements. So, the element beryllium for a long time was considered an analogue of aluminum and the formula was attributed to its oxide. Based on the percentage composition and the expected formula of beryllium oxide, its atomic mass was considered to be 13.5. The periodic table has shown that there is only one place for beryllium in the table, namely above magnesium, so its oxide must have the formula , which gives the atomic mass of beryllium equal to ten. This conclusion was soon confirmed by determinations of the atomic mass of beryllium from the vapor density of its chloride.
Exactly And at present, the periodic law remains the guiding thread and guiding principle of chemistry. It was on its basis that transuranium elements located in the periodic table after uranium were artificially created in recent decades. One of them - element No. 101, first obtained in 1955 - was named mendelevium in honor of the great Russian scientist.
The discovery of the periodic law and the creation of a system of chemical elements was of great importance not only for chemistry, but also for philosophy, for our entire understanding of the world. Mendeleev showed that chemical elements form a harmonious system, which is based on a fundamental law of nature. This is an expression of the position of materialist dialectics about the interconnection and interdependence of natural phenomena. Revealing the relationship between the properties of chemical elements and the mass of their atoms, the periodic law was a brilliant confirmation of one of the universal laws of the development of nature - the law of the transition of quantity into quality.
The subsequent development of science made it possible, based on the periodic law, to understand the structure of matter much more deeply than was possible during Mendeleev’s lifetime.
The theory of atomic structure developed in the 20th century, in turn, gave the periodic law and the periodic system of elements a new, deeper illumination. The prophetic words of Mendeleev were brilliantly confirmed: “The periodic law is not threatened with destruction, but only superstructure and development are promised.”
periodic law of mendeley atom
The periodic law made it possible to systematize and generalize a huge amount of scientific information in chemistry. This function of the law is usually called integrative. It is especially clearly manifested in the structuring of scientific and educational material chemistry. Academician A.E. Fersman said that the system united all chemistry within a single spatial, chronological, genetic, and energetic connection.
The integrative role of the Periodic Law was also manifested in the fact that some data on elements that allegedly fell out of general patterns, were checked and clarified by both the author himself and his followers.
This happened with the characteristics of beryllium. Before Mendeleev's work, it was considered a trivalent analogue of aluminum due to their so-called diagonal similarity. Thus, in the second period there were two trivalent elements and not a single divalent one. It was at this stage, first at the level of mental model constructions, that Mendeleev suspected an error in the studies of the properties of beryllium. He then found the work of Russian chemist Avdeev, who argued that beryllium was divalent and had an atomic weight of 9. Avdeev's work went unnoticed scientific world, the author died early, apparently having been poisoned by extremely poisonous beryllium compounds. The results of Avdeev’s research were established in science thanks to the Periodic Law.
Such changes and refinements of the values of both atomic weights and valences were made by Mendeleev for nine more elements (In, V, Th, U, La, Ce and three other lanthanides). For ten more elements, only atomic weights were corrected. And all these clarifications were subsequently confirmed experimentally.
In the same way, the work of Karl Karlovich Klaus helped Mendeleev form a unique VIII group of elements, explaining the horizontal and vertical similarities in the triads of elements:
iron cobalt nickel
ruthenium rhodium palladium
octagonal iridium platinum
The prognostic (predictive) function of the Periodic Law received the most striking confirmation in the discovery of unknown elements with serial numbers 21, 31 and 32. Their existence was first predicted on an intuitive level, but with the formation of the Mendeleev system with high degree I was able to accurately calculate their properties. Fine famous story The discovery of scandium, gallium and germanium was a triumph of Mendeleev's discovery. F. Engels wrote: “By unconsciously applying the Hegelian law of the transition of quantity into quality, Mendeleev accomplished a scientific feat that can safely be placed next to the discovery of Laverrier, who calculated the orbit of the unknown planet Neptune.” However, there is a desire to argue with the classic. Firstly, all of Mendeleev’s research, starting from his student years, was quite consciously based on Hegel’s law. Secondly, Laverrier calculated the orbit of Neptune according to Newton’s long-known and proven laws, and D.I. Mendeleev made all predictions on the basis of the universal law of nature discovered by himself.
At the end of his life, Mendeleev noted with satisfaction: “Having written in 1871 an article on the application of the periodic law to determining the properties of elements not yet discovered, I did not think that I would live to justify this consequence of the periodic law, but reality answered differently. I described three elements: ekaboron, ekaaluminum and ekasilicon, and less than 20 years later I had the greatest joy of seeing all three discovered... L. de Boisbaudran, Nilsson and Winkler, for my part, I consider true strengtheners of the periodic law. Without them, he would not have been recognized to the extent that he has now.” In total, Mendeleev predicted twelve elements.
From the very beginning, Mendeleev pointed out that the law describes the properties not only of the chemical elements themselves, but also of many of their compounds, including hitherto unknown ones. To confirm this, it is enough to give the following example. Since 1929, when Academician P. L. Kapitsa first discovered the non-metallic conductivity of germanium, the development of the study of semiconductors began in all countries of the world. It immediately became clear that elements with such properties occupy the main subgroup of group IV. Over time, the understanding came that semiconductor properties should, to a greater or lesser extent, be possessed by compounds of elements located in periods equally distant from this group (for example, with general formula type AzV;). This immediately made the search for new practically important semiconductors targeted and predictable. Almost all modern electronics are based on such connections.
It is important to note that predictions within the Periodic Table were made even after its general acceptance. In 1913 Moseley discovered that the wavelength of X-rays that are received from anticathodes made of different elements, changes naturally depending on the serial number conventionally assigned to the elements in the Periodic Table. The experiment confirmed that the serial number of an element has a direct physical meaning. Only later were serial numbers related to the value of the positive charge of the nucleus. But Moseley’s law made it possible to immediately experimentally confirm the number of elements in periods and at the same time predict the places of hafnium (No. 72) and rhenium (No. 75) that had not yet been discovered by that time.
The same studies by Moseley made it possible to remove the serious “headache” that certain deviations from the correct series of increasing atomic masses of elements in the table of atomic masses caused Mendeleev. Mendeleev made them under the pressure of chemical analogies, partly at an expert level, and partly simply at an intuitive level. For example, cobalt was ahead of nickel in the table, and iodine, with a lower atomic weight, followed the heavier tellurium. IN natural sciences It has long been known that one “ugly” fact that does not fit into the framework of the most beautiful theory can destroy it. Likewise, unexplained deviations threatened the Periodic Law. But Moseley experimentally proved that the serial numbers of cobalt (No. 27) and nickel (No. 28) exactly correspond to their position in the system. It turned out that these exceptions only confirm the general rule.
An important prediction was made in 1883 by Nikolai Aleksandrovich Morozov. For participation in the People's Will movement, chemistry student Morozov was sentenced to death, which was later replaced by life imprisonment in solitary confinement. He spent about thirty years in royal prisons. A prisoner of the Shlisselburg fortress had the opportunity to receive some scientific literature in chemistry. Based on the analysis of the intervals of atomic weights between neighboring groups elements in the periodic table Morozov came to the intuitive conclusion about the possibility of the existence of another group of unknown elements between the groups of halogens and alkali metals with “ null properties" He suggested looking for them in the air. Moreover, he expressed a hypothesis about the structure of atoms and, on its basis, tried to reveal the causes of periodicity in the properties of elements.
However, Morozov's hypotheses became available for discussion much later, when he was released after the events of 1905. But by that time, inert gases had already been discovered and studied.
For a long time, the fact of the existence of inert gases and their position in the periodic table caused serious controversy in the chemical world. Mendeleev himself believed for some time that an unknown simple substance of the Nj type could be hiding under the brand name of open argon. The first rational assumption about the place of inert gases was made by the author of their discovery, William Ramsay. And in 1906, Mendeleev wrote: “When the Periodic Table was established (18b9), not only was argon not known, but there was no reason to suspect the possibility of the existence of such elements. Today... these elements, in terms of their atomic weights, have taken the exact place between the halogens and the alkali metals.”
For a long time there was a debate: to allocate inert gases into an independent zero group of elements or to consider them as the main subgroup of group VIII. Each point of view has its pros and cons.
Based on the position of the elements in the Periodic Table, theoretical chemists led by Linus Pauling have long doubted the complete chemical passivity of noble gases, directly pointing to the possible stability of their fluorides and oxides. But only in 1962, the American chemist Neil Bartlett was the first to carry out the reaction of platinum hexafluoride with oxygen under the most ordinary conditions, obtaining xenon hexafluoroplatinate XePtF^, followed by other gas compounds, which are now more correctly called noble rather than inert.
The periodic law retains its predictive function to this day.
It should be noted that predictions of unknown members of any set can be of two types. If the properties of an element located within a known series of similar ones are predicted, then such a prediction is called interpolation. It is natural to assume that these properties will be subject to the same laws as the properties of neighboring elements. This is how the properties of the missing elements within the periodic table were predicted. It is much more difficult to predict the characteristics of new members of sets if they are outside the described part. Extrapolation - the prediction of function values that are outside a number of known patterns - is always less certain.
It was this problem that confronted scientists when they began searching for elements beyond the known boundaries of the system. At the beginning of the 20th century. The periodic table ended with uranium (No. 92). The first attempts to obtain transuranium elements were made in 1934, when Enrico Fermi and Emilio Segre bombarded uranium with neutrons. Thus began the road to actinoids and transactinoids.
Nuclear reactions are also used to synthesize other previously unknown elements.
Element No. 101, artificially synthesized by Eienn Theodor Seaborg and his colleagues, was named “mendelevium”. Seaborg himself said this: “It is especially significant to note that element 101 was named in honor of the great Russian chemist D.I. Mendeleev by American scientists, who always considered him a pioneer in chemistry.”
The number of newly discovered, or rather artificially created, elements is constantly growing. Synthesis of the most heavy nuclei elements with serial numbers 113 and 115 were carried out at the Russian Joint Institute nuclear research in Dubna by bombarding nuclei of artificially obtained americium with nuclei of the heavy isotope calcium-48. In this case, the nucleus of element No. 115 appears, which immediately decays to form the nucleus of element No. 113. Such superheavy elements do not exist in nature, but they arise during supernova explosions, and could also exist during the period big bang. Their research helps to understand how our Universe came into being.
A total of 39 naturally occurring radioactive isotopes occur in nature. Different isotopes decay at different rates, which are characterized by half-lives. The half-life of uranium-238 is 4.5 billion years, and for some other elements it can be equal to millionths of a second.
Radioactive elements, sequentially decaying and transforming into each other, form entire series. Three such series are known: according to the initial element, all members of the series are combined into the families of uranium, actinouranium and thorium. Another family consists of artificially produced radioactive isotopes. In all families, the transformations are completed by the appearance of non-radioactive lead atoms.
Since in earth's crust can only contain isotopes whose half-lives are commensurate with the age of the Earth, it can be assumed that over billions of years of its history there also existed such short-lived isotopes that by now literally this word has died out. These probably included the heavy isotope potassium-40. As a result of its complete decay, the tabulated value of the atomic mass of potassium today is 39.102, so it is inferior in mass to element No. 18 argon (39.948). This explains the exceptions in the consistent increase in atomic masses of elements in the periodic table.
Academician V. I. Goldansky, in a speech dedicated to the memory of Mendeleev, noted “the fundamental role that Mendeleev’s works play even in completely new areas of chemistry, which arose decades after the death of the brilliant creator of the Periodic Table.”
Science is the history and repository of the wisdom and experience of centuries, their rational contemplation and tested judgment.
D. I. Mendeleev
It rarely happens that scientific discovery turned out to be something completely unexpected; it is almost always anticipated:
However, subsequent generations, who use proven answers to all questions, often find it difficult to appreciate what difficulties it cost their predecessors.
C. Darwin
Each of the sciences about the world around us has as its subject of study specific forms of movement of matter. The prevailing ideas consider these forms of movement in order of increasing complexity:
mechanical - physical - chemical - biological - social. Each of the subsequent forms does not reject the previous ones, but includes them.
It is no coincidence that at the celebration of the centenary of the discovery of the Periodic Law, G. T. Seaborg devoted his report to the latest achievements of chemistry. In it, he highly appreciated the amazing achievements of the Russian scientist: “When considering the evolution of the Periodic Table since the time of Mendeleev, the most striking thing is that he was able to create the Periodic Table of elements, although Mendeleev was not aware of such now generally accepted concepts as nuclear structure and isotopes , relationship between atomic numbers and valence, electronic nature of atoms, periodicity of chemical properties, explained electronic structure, and, finally, radioactivity."
One can cite the words of Academician A.E. Fersman, who drew attention to the future: “New theories, brilliant generalizations will appear and die. New ideas will replace our already outdated concepts of the atom and electron. Greatest discoveries and experiments will nullify the past and open today horizons of incredible novelty and breadth - all this will come and go, but Mendeleev’s Periodic Law will always live and guide the search.”
The periodic table of elements has become one of the most valuable generalizations in chemistry. It is like a summary of the chemistry of all elements, a graph from which you can read the properties of the elements and their compounds. The system made it possible to clarify the position, atomic masses, and valency values of some elements. Based on the table, it was possible to predict the existence and properties of yet undiscovered elements. Mendeleev formulated the periodic law and proposed its graphical representation, but at that time it was impossible to determine the nature of periodicity. The meaning of the periodic law was revealed later, in connection with discoveries on the structure of the atom.
1. In what year was the periodic law discovered?
2. What did Mendeleev take as the basis for the systematization of elements?
3. What does the law discovered by Mendeleev say?
4. What is the difference with the modern formulation?
5. What is called an atomic orbital?
6. How do properties change over periods?
7. How are periods divided?
8. What is a group called?
9. How are groups divided?
10. What types of electrons do you know?
11. How filling occurs energy levels?
Lecture No. 4: Valency and oxidation state. Frequency of property changes.
Origin of the concept of valence. The valence of chemical elements is one of their most important properties. The concept of valence was introduced into science by E. Frankland in 1852. At first, the concept was exclusively stoichiometric in nature and stemmed from the law of equivalents. The meaning of the concept of valence follows from a comparison of the values of atomic mass and the equivalent of chemical elements.
With the establishment of atomic-molecular concepts, the concept of valence acquired a certain structural and theoretical meaning. Valency began to be understood as the ability of one atom of a given element to attach to itself a certain number of atoms of another chemical element. The corresponding capacity of the hydrogen atom was taken as the unit of valence, since the ratio of the atomic mass of hydrogen to its equivalent is equal to unity. Thus, the valency of a chemical element was defined as the ability of its atom to attach a certain number of hydrogen atoms. If a given element did not form compounds with hydrogen, its valence was determined as the ability of its atom to replace a certain number of hydrogen atoms in its compounds.
This idea of valence was confirmed for the simplest compounds.
Based on the idea of the valence of elements, the idea of the valence of entire groups arose. So, for example, the OH group, since it added one hydrogen atom or replaced one hydrogen atom in its other compounds, was assigned a valence of one. However, the idea of valency lost its unambiguity when it came to more complex compounds. So, for example, in hydrogen peroxide H 2 O 2 the valency of oxygen should be recognized as equal to one, since in this compound there is one hydrogen atom for each oxygen atom. However, it is known that each oxygen atom in H 2 O 2 is connected to one hydrogen atom and one monovalent OH group, i.e. oxygen is divalent. Similarly, the valency of carbon in ethane C 2 H 6 should be recognized as equal to three, since in this compound there are three hydrogen atoms for each carbon atom, but since each carbon atom is connected to three hydrogen atoms and one monovalent group CH 3, the valence carbon in C 2 H 6 is equal to four.
It should be noted that when forming ideas about the valence of individual elements, these complicating circumstances were not taken into account, and only the composition of the simplest compounds was taken into account. But even at the same time, it turned out that for many elements the valency in different compounds is not the same. This was especially noticeable for compounds of some elements with hydrogen and oxygen, in which different valences appeared. Thus, in combination with hydrogen, the valency of sulfur turned out to be equal to two, and with oxygen – six. Therefore, they began to distinguish between valency for hydrogen and valence for oxygen.
Subsequently, in connection with the idea that in compounds some atoms are polarized positively and others negatively, the concept of valence in oxygen and hydrogen compounds was replaced by the concept of positive and negative valency.
Various meanings The valencies of the same elements were also manifested in their various compounds with oxygen. In other words, the same elements were able to exhibit different positive valence. This is how the idea of variable positive valence of some elements appeared. As for the negative valence of non-metallic elements, it, as a rule, turned out to be constant for the same elements.
The majority of elements exhibited variable positive valence. However, each of these elements was characterized by its maximum valency. This maximum valency is called characteristic.
Later, in connection with the emergence and development of the electronic theory of atomic structure and chemical bonds, valence began to be associated with the number of electrons passing from one atom to another, or with the number of chemical bonds arising between atoms in the process of formation chemical compound.
Electrovalency and covalency. The positive or negative valence of an element is most easily determined if two elements formed an ionic compound: the element whose atom became a positively charged ion was considered to have a positive valence, and the element whose atom became a negatively charged ion had a negative valence. The numerical value of valence was considered equal to the magnitude of the ion charge. Since ions in compounds are formed by the donation and acquisition of electrons by atoms, the amount of charge of the ions is determined by the number of electrons given up (positive) and added (negative) by the atoms. In accordance with this, the positive valence of an element was measured by the number of electrons donated by its atom, and the negative valence - by the number of electrons attached by a given atom. Thus, since valence was measured by the magnitude of the electric charge of atoms, it received the name electrovalency. It is also called ionic valency.
Among chemical compounds there are those in whose molecules the atoms are not polarized. Obviously, for them the concept of positive and negative electrovalency is not applicable. If the molecule is composed of atoms of one element (elementary substances), the usual concept of stoichiometric valency loses its meaning. However, in order to evaluate the ability of atoms to attach a given number of other atoms, they began to use the number of chemical bonds that arise between a given atom and other atoms during the formation of a chemical compound. Because these chemical bonds, which are electron pairs that simultaneously belong to both connected atoms, are called covalent; the ability of an atom to form one or another number of chemical bonds with other atoms is called covalency. To establish covalency, structural formulas are used in which chemical bonds are represented by dashes.
Oxidation state and oxidation number. In reactions of the formation of ionic compounds, the transition of electrons from one reacting atoms or ions to others is accompanied by a corresponding change in the value or sign of their electrovalence. When compounds of a covalent nature are formed, such a change in the electrovalent state of the atoms actually does not occur, but only a redistribution of electronic bonds takes place, and the valence of the original reacting substances does not change. Currently, to characterize the state of an element in connections, a conditional concept has been introduced oxidation states. The numerical expression of the oxidation state is called oxidation number.
The oxidation numbers of atoms can have positive, zero and negative values. A positive oxidation number is determined by the number of electrons drawn from a given atom, and a negative oxidation number is determined by the number of electrons attracted by a given atom. The oxidation number can be assigned to each atom in any substance, for which one must be guided by the following simple rules:
1. The oxidation numbers of atoms in any elementary substances are zero.
2. The oxidation numbers of elementary ions in substances of ionic nature are equal to the values electric charges these ions.
3. The oxidation numbers of atoms in compounds of a covalent nature are determined by the conventional calculation that each electron drawn from an atom gives it a charge equal to +1, and each electron attracted gives it a charge equal to –1.
4. The algebraic sum of the oxidation numbers of all atoms of any compound is zero.
5. The fluorine atom in all its compounds with other elements has an oxidation number of –1.
The determination of the oxidation state is associated with the concept of electronegativity of elements. Using this concept, another rule is formulated.
6. In compounds, the oxidation number is negative for atoms of elements with higher electronegativity and positive for atoms of elements with lower electronegativity.
The concept of oxidation state has thus replaced the concept of electrovalence. In this regard, it seems inappropriate to use the concept of covalency. To characterize elements, it is better to use the concept of valence, defining it by the number of electrons used by a given atom to form electron pairs, regardless of whether they are attracted to a given atom or, conversely, withdrawn from it. Then the valency will be expressed as an unsigned number. In contrast to valence, the oxidation state is determined by the number of electrons drawn from a given atom (positive), or attracted to it (negative). In many cases, the arithmetic values of valency and oxidation state coincide - this is quite natural. In some cases, the numerical values of valency and oxidation state differ from each other. For example, in molecules of free halogens the valence of both atoms is equal to one, and the oxidation state is zero. In the molecules of oxygen and hydrogen peroxide, the valency of both oxygen atoms is two, and their oxidation state in the oxygen molecule is zero, and in the hydrogen peroxide molecule it is minus one. In the molecules of nitrogen and hydrazine - N 4 H 2 - the valency of both nitrogen atoms is three, and the oxidation state in the elemental nitrogen molecule is zero, and in the hydrazine molecule it is minus two.
It is obvious that valence characterizes atoms that are only part of any compound, even a homonuclear one, that is, consisting of atoms of one element; It makes no sense to talk about the valency of individual atoms. The degree of oxidation characterizes the state of atoms both included in a compound and existing separately.
Questions to reinforce the topic:
1. Who introduced the concept of “valence”?
2. What is valency called?
3. What is the difference between valency and oxidation state?
4. What is the valency?
5. How is the oxidation state determined?
6. Are the valency and oxidation state of an element always equal?
7. By which element is the valence of an element determined?
8. What characterizes the valence of an element, and what is the oxidation state?
9. Can the valency of an element be negative?
Lecture No. 5: The rate of a chemical reaction.
Chemical reactions may vary significantly in duration. A mixture of hydrogen and oxygen at room temperature can remain virtually unchanged for a long time, but if struck or ignited, an explosion will occur. The iron plate slowly rusts, and a piece of white phosphorus spontaneously ignites in air. It is important to know how quickly a particular reaction occurs in order to be able to control its progress.
Scientific significance of the periodic law. Life and work of D.I. Mendeleev
Discovery of the periodic law and creation of the Periodic Table of Chemical Elements - greatest achievement science of the 19th century. Experimental confirmation of the relative atomic masses changed by D.I. Mendeleev, the discovery of elements with the properties envisaged by him, and the location of open inert gases in the periodic table led to the universal recognition of the periodic law.
The discovery of the periodic law led to further rapid development of chemistry: over the next thirty years, 20 new chemical elements were discovered. The periodic law contributed further development works on the study of the structure of the atom, as a result of which the relationship between the structure of the atom and the periodic change of their properties was established. Based on the periodic law, scientists were able to extract substances with given properties and synthesize new chemical elements. The periodic law has allowed scientists to build hypotheses about the evolution of chemical elements in the Universe.
The periodic law of D.I. Mendeleev has general scientific significance and is a fundamental law of nature.
Dmitry Ivanovich Mendeleev was born in 1834 in Tobolsk. After graduating from the Tobolsk gymnasium, he studied at the St. Petersburg Pedagogical Institute, from which he graduated with a gold medal. As a student, D.I. Mendeleev began studying scientific research. After studying, he spent two years abroad in the laboratory of the famous chemist Robert Bunsen. In 1863, he was elected professor, first at the St. Petersburg Institute of Technology, and subsequently at St. Petersburg University.
Mendeleev conducted research in the field chemical nature solutions, state of gases, heat of combustion of fuel. He was interested various problems Agriculture, mining, metallurgy issues, worked on the problem of underground gasification of fuel, studied petroleum engineering. The most significant result creative activity, brought D.I. Mendeleev worldwide fame, was the discovery in 1869 of the Periodic Law and the Periodic Table of Chemical Elements. He wrote about 500 articles on chemistry, physics, technology, economics, and geodesy. He organized and was the director of the first Russian Chamber of Weights and Measures, and concluded the beginning of modern metrology. Invented general equation state of an ideal gas, generalized the Clapeyron equation (Clapeyron-Mendeleev equation).
Mendeleev lived to be 73 years old. For his achievements, he was elected a member of 90 foreign academies of sciences and honorary doctorates of many universities. The 101st chemical element (Mendelevium) is named in his honor.
