Periodic system of chemical elements. Chapter ii. Atomic structure and periodic law
The periodic law is the basis of modern chemistry. The knowledge of the periodic law is the basis for all scientific areas and research in chemistry: the study of the mutual transformations of substances, the production of new materials, the theoretical study of the structure of substances, types of chemical bonds, and so on.
The nuclear charge determines the number of electrons in an atom, each subsequent element has one electron more than the previous one. The nuclear charge determines the structure of the electron shell of the atom in the ground state. Elements are arranged in a periodic system of elements in order of increasing charge of the nuclei of their atoms. The elements the electronic configurations of atoms are periodically repeated and, as a consequence of this, the chemical properties that are determined are periodically repeated electron configuration of atoms. The periodicity of the electronic structure is manifested in the fact that s-, p-, and d-elements with identical configurations of electronic sublevels are repeated again after a certain number of elements. The periodicity is inherent in the entire electron shell of atoms, and not just its outer layers. The periodicity of electronic structures leads to a periodic change in a number of chemical and physical properties of elements: atomic radii, ionization energies, electron affinity, and electronegativity. We will discuss this more specifically.
Atomic Radii chemical elements change periodically depending on the charge of the atomic nucleus (or the serial number of the element). In periods, the radii of atoms decrease from alkali metal to halogen. So the atomic radius of the sodium atom is 0.186 nm, magnesium is 0.16 nm, and chlorine is 0.099 nm. The atomic radius of the next alkali metal, which opens the subsequent period, increases sharply, its radius is much larger than the radius of the alkali metal standing above it. For example: the radius of the sodium atom is 0.186 nm, and the potassium atom is 0.231 nm.
The decrease in the radii of atoms in periods from left to right, that is, with an increase in the charge of the atomic nucleus, is explained by the fact that an increase in the charge of the atomic nucleus contributes to a stronger attraction of the electrons of a given electronic level to the nucleus (it acts stronger than the repulsion of electrons from each other).
In groups, with increasing charge of the atomic nucleus (from top to bottom), the radii of atoms increase. This is due to the fact that each element below has one electronic level more, so it also has an atomic radius. This pattern is more pronounced in elements of the main subgroups (in s- and p-elements) than in elements of secondary subgroups (d-elements).
There are exceptions to these considered patterns, but we will not discuss them, since this is not part of our program.
We also point out that it is necessary to distinguish between the radii of a free atom and the following radii:
but) covalent radius is half the internuclear distance in the molecules or crystals of the corresponding simple substances (i.e., substances with a covalent bond type);
b) the metal radius is half the distance between the centers of two neighboring atoms in the crystal lattice of the metal;
in) the ionic radii of atoms are considered as half the distance of the sum of the radii of the cation and anion (it should be remembered that the radii of cations are always smaller than the atomic radii of the corresponding elements, and the radii of anions are larger than the radii of the atoms of the corresponding elements).
Ionization energy and electron affinity these are parameters that allow us to evaluate the ability of atoms to lose and receive electrons.
The periodic law and the periodic system of chemical elements D. I. Mendeleev based on ideas about the structure of atoms. The importance of the periodic law for the development of science.
In 1869, D. I. Mendeleev, based on an analysis of the properties of simple substances and compounds, formulated the Periodic Law:
The properties of simple bodies ... and compounds of elements are periodically dependent on the atomic masses of the elements.
Based on the periodic law, a periodic system of elements was compiled. In it, elements with similar properties were combined into vertical columns - groups. In some cases, when placing elements in the Periodic system, it was necessary to disrupt the sequence of increase in atomic masses in order to observe the periodicity of repetition of properties. For example, I had to "swap" tellurium and iodine, as well as argon and potassium.
The reason is that Mendeleev proposed a periodic law at a time when nothing was known about the structure of the atom.
After the planetary model of the atom was proposed in the 20th century, the periodic law is formulated as follows:
The properties of chemical elements and compounds are periodically dependent on the charges of atomic nuclei.
The nuclear charge is equal to the number of the element in the periodic system and the number of electrons in the electron shell of the atom.
This wording explained the "violations" of the Periodic Law.
In the Periodic system, the period number is equal to the number of electronic levels in the atom, the group number for elements of the main subgroups is equal to the number of electrons at the external level.
The reason for the periodic change in the properties of chemical elements is the periodic filling of electronic shells. After filling in the next shell, a new period begins. Periodic changes in elements are clearly visible on changes in the composition and properties and properties of oxides.
The scientific significance of the periodic law. The periodic law made it possible to systematize the properties of chemical elements and their compounds. In compiling the periodic system, Mendeleev predicted the existence of many elements not yet discovered, leaving free cells for them, and predicted many properties of undiscovered elements, which facilitated their discovery
Ticket number 2
The structure of atoms of chemical elements on the example of elements of the second period and IV-A group of the periodic system of chemical elements D. I. Mendeleev. Patterns in changing the properties of these chemical elements and the simple and complex substances (oxides, hydroxides) formed by them, depending on the structure of their atoms.
When moving from left to right along the period, the metal properties of the elements become less pronounced. When moving from top to bottom within the same group, elements, on the contrary, exhibit increasingly pronounced metallic properties. Elements located in the middle of short periods (2nd and 3rd periods), as a rule, have a frame covalent structure, and elements from the right side of these periods exist in the form of simple covalent molecules.
The atomic radii change as follows: decrease when moving from left to right along the period; increase when moving from top to bottom along the group. When moving from left to right over a period, electronegativity, ionization energy and electron affinity increase, which reach a maximum for halogens. In noble gases, the electronegativity is equal to 0. The change in the electron affinity of the elements when moving from top to bottom along the group is not so characteristic, but the electronegativity of the elements decreases.
In the elements of the second period, 2s are filled, and then 2p orbitals.
The main subgroup of group IV of the periodic system of chemical elements D. M. Mendeleev contains carbon C, silicon Si, germanium Ge, tin Sn and lead Pb. The outer electronic layer of these elements contains 4 electrons (configuration s 2 p 2). Therefore, the elements of the carbon subgroup must have some similarities. In particular, their highest oxidation state is the same and equal to +4.
And what is the reason for the difference in the properties of the elements of the subgroup? The difference between the ionization energy and the radius of their atoms. With an increase in the atomic number, the properties of elements naturally change. So, carbon and silicon are typical non-metals, tin and lead are metals. This is manifested primarily in the fact that carbon forms a simple non-metal substance (diamond), and lead is a typical metal.
Germanium is intermediate. According to the structure of the electron shell of an atom, p-elements of group IV have even oxidation states: +4, +2, - 4. The formula of the simplest hydrogen compounds is EN 4, and the E – H bonds are covalent and equivalent due to the hybridization of s and p orbitals to form directed at tetrahedral angles of sp 3 orbitals.
The weakening of the signs of a nonmetallic element means that in the subgroup (С-Si-Ge-Sn-Pb), the highest positive oxidation state +4 is becoming less and less characteristic, and the oxidation state +2 is becoming more typical. So, if for carbon the most stable compounds in which it has an oxidation state of +4, then for lead stable compounds in which it exhibits an oxidation state of +2.
And what about the stability of compounds of elements in the negative oxidation state of -4? Compared with non-metallic elements of the VII-V groups, the signs of the non-metallic element of the p-elements of group IV are less pronounced. Therefore, for elements of the carbon subgroup, a negative oxidation state is not typical.
The periodic system of chemical elements is a classification of chemical elements created by D. I. Mendeleev on the basis of the periodic law discovered by him in 1869.
D.I. Mendeleev
According to the modern formulation of this law, in a continuous series of elements arranged in increasing order of the positive charge of the nuclei of their atoms, elements with similar properties are periodically repeated.
The periodic system of chemical elements, presented in the form of a table, consists of periods, series and groups.
At the beginning of each period (with the exception of the first) there is an element with pronounced metallic properties (alkali metal).
Symbols for the color table: 1 - chemical sign of the element; 2 - name; 3 - atomic mass (atomic weight); 4 - serial number; 5 - electron distribution in layers.
As the element serial number increases, which is equal to the positive charge of the nucleus of its atom, the metallic properties gradually weaken and non-metallic properties increase. The penultimate element in each period is an element with pronounced non-metallic properties (), and the last is an inert gas. In the I period there are 2 elements, in II and III - 8 elements each, in IV and V - 18 each, in VI - 32 and in VII (incomplete period) - 17 elements.
The first three periods are called small periods, each of them consists of one horizontal row; the rest are in large periods, each of which (except for the VII period) consists of two horizontal rows - even (upper) and odd (lower). In even rows of large periods are only metals. The properties of the elements in these rows change weakly with an increasing number. The properties of elements in the odd rows of large periods change. In the VI period, 14 elements, very similar in chemical properties, follow lanthanum. These elements, called lanthanides, are listed separately below the main table. Actinides, the elements following actinium, are presented in the table in the same way.
The table has nine vertical groups. The group number, with rare exceptions, is equal to the highest positive valency of the elements of this group. Each group, except zero and eighth, is divided into subgroups. - main (located to the right) and secondary. In the main subgroups, with an increase in the serial number, metallic and non-metallic properties of elements intensify.
Thus, the chemical and a number of physical properties of elements are determined by the place that this element occupies in the periodic system.
Biogenic elements, i.e., elements that make up organisms and fulfill a certain biological role in it, occupy the top of the periodic table. The cells occupied by the elements that make up the bulk (more than 99%) of living matter are painted blue, the cells occupied by trace elements are pink (see).
The periodic system of chemical elements is the greatest achievement of modern science and a vivid expression of the most general dialectical laws of nature.
See also Atomic Weight.
The periodic system of chemical elements is a natural classification of chemical elements created by D. I. Mendeleev on the basis of the periodic law discovered by him in 1869.
In the original formulation, the periodic law of D. I. Mendeleev stated: the properties of chemical elements, as well as the forms and properties of their compounds, are periodically dependent on the atomic weights of the elements. Later, with the development of the doctrine of the structure of the atom, it was shown that a more accurate characteristic of each element is not atomic weight (see), but the value of the positive charge of the atomic nucleus of an element, equal to the ordinal (atomic) number of this element in the periodic system of D. I. Mendeleev . The number of positive charges of the nucleus of an atom is equal to the number of electrons surrounding the nucleus of an atom, since atoms are generally electrically neutral. In the light of these data, the periodic law is formulated as follows: the properties of chemical elements, as well as the forms and properties of their compounds, are periodically dependent on the magnitude of the positive charge of the nuclei of their atoms. This means that in a continuous series of elements arranged in increasing order of the positive charges of the nuclei of their atoms, elements with similar properties will periodically repeat.
The tabular form of the periodic system of chemical elements is presented in its modern form. It consists of periods, series and groups. The period represents a consecutive horizontal row of elements arranged in increasing order of the positive charge of the nuclei of their atoms.
At the beginning of each period (with the exception of the first) there is an element with pronounced metallic properties (alkali metal). Then, as the serial number increases, the metal properties gradually weaken and the non-metallic properties of the elements increase. The penultimate element in each period is an element with pronounced non-metallic properties (halogen), and the last is an inert gas. I period consists of two elements, the role of alkali metal and halogen here is simultaneously performed by hydrogen. II and III periods include 8 elements, called Mendeleev typical. IV and V periods consist of 18 elements, VI-32. VII period has not yet been completed and replenished with artificially created elements; there are currently 17 elements in this period. Periods I, II and III are called small, each of them consists of one horizontal row, IV-VII - large: they (with the exception of VII) include two horizontal rows - even (upper) and odd (lower). Only even metals are in even rows of large periods, and the change in the properties of elements in a row from left to right is weakly expressed.
In odd rows of large periods, the properties of the elements in the series change in the same way as the properties of typical elements. In an even row of the VI period, after lanthanum there are 14 elements [called lanthanides (see), lanthanides, rare earth elements], similar in chemical properties to lanthanum and to each other. Their list is given separately under the table.
The elements following the actinium actinides (actinides) are separately written out and shown below the table.
In the periodic system of chemical elements, nine groups are located vertically. The group number is equal to the highest positive valency (see) of the elements of this group. The exceptions are fluorine (it can only be negatively monovalent) and bromine (it is not heptavalent); in addition, copper, silver, gold can exhibit a valency of more than +1 (Cu-1 and 2, Ag and Au-1 and 3), and of the elements of group VIII, only osmium and ruthenium have a valency of +8. Each group, except for the eighth and zero, is divided into two subgroups: the main (located to the right) and the side. The main subgroups include typical elements and elements of large periods, the side - only elements of large periods and, moreover, metals.
By chemical properties, the elements of each subgroup of this group are significantly different from each other and only the highest positive valency is the same for all elements of this group. In the main subgroups, the metallic properties of the elements are enhanced from top to bottom and non-metallic are weakened (for example, francium is the element with the most pronounced metallic properties, and fluorine is non-metallic). Thus, the place of an element in the periodic table (serial number) determines its properties, which are the average of the properties of neighboring elements vertically and horizontally.
Some groups of elements have special names. So, the elements of the main subgroups of group I are called alkali metals, group II - alkaline earth metals, group VII - halogens, elements located behind uranium - transuranic. Elements that are part of organisms participate in metabolic processes and have a pronounced biological role, called nutrients. All of them occupy the top of the table of D. I. Mendeleev. This is primarily O, C, H, N, Ca, P, K, S, Na, Cl, Mg and Fe, which make up the bulk of living matter (more than 99%). The places occupied by these elements in the periodic system are painted in light blue. Biogenic elements, which are very few in the body (from 10 -3 to 10 -14%), are called microelements (see). In the cells of the periodic system, stained yellow, microelements are placed, the vital importance of which is proved for humans.
According to the theory of the structure of atoms (see Atom), the chemical properties of elements depend mainly on the number of electrons on the outer electron shell. The periodic change in the properties of elements with an increase in the positive charge of atomic nuclei is explained by the periodic repetition of the structure of the outer electron shell (energy level) of atoms.
In small periods, with an increase in the positive charge of the nucleus, the number of electrons on the outer shell increases from 1 to 2 in the I period and from 1 to 8 in the II and III periods. Hence the change in the properties of elements in the period from an alkali metal to an inert gas. The outer electron shell containing 8 electrons is complete and energetically stable (elements of the zero group are chemically inert).
In large periods in even rows, with an increase in the positive charge of the nuclei, the number of electrons on the outer shell remains constant (1 or 2) and electrons are filling in the second outside the shell. Hence the slow change in the properties of elements in even rows. In odd rows of large periods with increasing nuclear charge, the outer shell is filled with electrons (from 1 to 8) and the properties of the elements change as in typical elements.
The number of electron shells in an atom is equal to the number of the period. The atoms of the elements of the main subgroups have on their outer shells the number of electrons equal to the group number. Atoms of elements of secondary subgroups contain one or two electrons on the outer shells. This explains the difference in the properties of the elements of the main and secondary subgroups. The group number indicates the possible number of electrons that can participate in the formation of chemical (valence) bonds (see. Molecule), therefore, such electrons are called valence. In the elements of secondary subgroups, not only the electrons of the outer shells, but also the penultimate ones are valence. The number and structure of electronic shells is indicated in the attached periodic system of chemical elements.
The periodic law of D. I. Mendeleev and the system based on it are of extremely great importance in science and practice. The periodic law and the system were the basis for the discovery of new chemical elements, the exact determination of their atomic weights, the development of the doctrine of the structure of atoms, the establishment of geochemical laws of the distribution of elements in the earth's crust, and the development of modern ideas about living matter, the composition of which and related patterns are in accordance with a periodic system. The biological activity of elements and their content in the body are also largely determined by the place that they occupy in the periodic table. So, with an increase in the serial number in a number of groups, the toxicity of elements increases and their content in the body decreases. The periodic law is a vivid expression of the most general dialectical laws of the development of nature.
Periodic Law D.I. Mendeleev.
The properties of chemical elements, and therefore the properties of simple and complex bodies formed by them, are periodically dependent on the value of atomic weight.
The physical meaning of the periodic law.
The physical meaning of the periodic law is to periodically change the properties of elements, as a result of periodically repeating e-shells of atoms, with a sequential increase in n.
The modern wording of PZ D.I. Mendeleev.
The property of chemical elements, as well as the property of simple or complex substances formed by them, is periodically dependent on the magnitude of the charge of the nuclei of their atoms.
Periodic system of elements.
Periodic system - a system of classifications of chemical elements created on the basis of a periodic law. Periodic system - establishes relationships between chemical elements reflecting their similarities and differences.
Periodic table (there are two types: short and long) of the elements.
Periodic table of elements - a graphical display of the periodic system of elements, consists of 7 periods and 8 groups.
Question 10
The periodic system and structure of the electron shells of atoms of elements.
It was further established that not only the serial number of the element has a deep physical meaning, but other concepts previously discussed also gradually acquired a physical meaning. For example, the group number, indicating the higher valency of an element, thereby reveals the maximum number of electrons of an atom of an element, which can participate in the formation of a chemical bond.
The period number, in turn, turned out to be related to the number of energy levels available in the electron shell of an atom of an element of a given period.
Thus, for example, the “coordinates” of tin Sn (serial number 50, period 5, the main subgroup of group IV) mean that there are 50 electrons in the tin atom, they are distributed at 5 energy levels, only 4 electrons are valence.
The physical meaning of finding elements in subgroups of various categories is extremely important. It turns out that for elements located in subgroups of category I, the next (last) electron is located on s-sublevel external level. These elements belong to the electronic family. For atoms of elements located in subgroups of category II, the next electron is located on sub-level external level. These are the elements of the “r” electronic family. Thus, the next 50th electron of tin atoms is located at the p-sublevel of the external, that is, the 5th energy level.
At atoms of elements of subgroups of category III, the next electron is located on d-sublevel, but already before the outer level, these are elements of the electronic family “d”. In atoms of lanthanides and actinides, the next electron is located on the f-sublevel, before the external level. These are elements of the electronic family. "F".
It is no accident, therefore, that the numbers of subgroups of these 4 categories noted above, that is, 2-6-10-14, coincide with the maximum numbers of electrons at the sublevels s-p-d-f.
But it turns out that it is possible to solve the problem of the order of filling the electron shell and derive the electronic formula for the atom of any element and based on a periodic system, which with sufficient clarity indicates the level and sublevel of each next electron. The periodic system also indicates the placement of elements one after another over periods, groups, subgroups and the distribution of their electrons by levels and sublevels, because each element has its own, which characterizes its last electron. As an example, let us examine the preparation of an electronic formula for an atom of an element of zirconium (Zr). The periodic system gives the indicators and the “coordinates” of this element: serial number 40, period 5, group IV, side subgroup. First conclusions: a) of all 40 electrons, b) these 40 electrons are distributed at five energy levels; 4 are valence, d) the next 40th electron entered the d-sublevel before the outer, that is, the fourth energy level .. Similar conclusions can be drawn about each of the 39 elements preceding zirconium, only the indicators and coordinates will be different each time.
Periodic Law D.I. Mendeleev is a fundamental law establishing a periodic change in the properties of chemical elements depending on an increase in the charges of the nuclei of their atoms. Opened D.I. Mendeleev in March 1869 when comparing the properties of all the elements known at that time and their atomic masses. Mendeleev first used the term “periodic law” in November 1870, and in October 1871 gave the final formulation of the Periodic Law: “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 that they form, are periodically dependent from their atomic weight. "
Hund’s rule: atomic orbitals belonging to one sublevel are each filled first with one electron, and then they are filled with second electrons.
The Hund rule is also called the principle of maximum multiplicity, i.e. the maximum possible parallel direction of the spins of the electrons of one energy sublevel.
At the highest energy level of a free atom, there can be no more than eight electrons.
Electrons located at the highest energy level of an atom (in the outer electron layer) are called external; the number of external electrons in an atom of any element is never more than eight. For many elements, it is the number of external electrons (with filled internal sublevels) that largely determines their chemical properties. For other electrons whose atoms have an unfilled internal sublevel, for example 3 d-sublevel of atoms of elements such as Sc, Ti, Cr, Mn, etc., chemical properties depend on the number of both internal and external electrons. All these electrons are called valence; in the abbreviated electronic formulas of atoms, they are written after the symbol of the atomic core, i.e. after expression in square brackets.
2.3. Periodic Law and the Periodic System of Elements
At the beginning of the 20th century, with the discovery of the atomic structure, it was found that the periodicity of changes in the properties of elements is determined not by atomic weight, but by a nuclear chargeequal to the atomic number and the number of electrons whose distribution over the electron shells of the atom of the element determines its chemical properties.
The modern wording of the Periodic Law states:
Further development of the periodic system is associated with filling in the empty cells of the table, in which more and more new elements were placed: noble gases, natural and artificially produced radioactive elements. In 2010, with the synthesis of 117 elements, the seventh period of the periodic system was completed. However, the problem of the lower boundary of the periodic table remains one of the most important in modern theoretical chemistry
The graphic (tabular) expression of the periodic law is the periodic system of elements developed by Mendeleev .
3 forms of the periodic table are more common than others: “short” (short-period), “long” (long-period), “extra-long”.
In the “super-long” version, each period occupies exactly one line. In the “long” version, lanthanides and actinides are taken out of the general table, making it more compact. In the “short” form of recording, in addition to this, the fourth and subsequent periods occupy 2 lines.
Elements arranged in ascending order of Z (H, He, Li, Be ...) form seven periods.
In periodsthe properties of elements naturally change during the transition from alkali metals to noble gases
