Periodic law of D.I. Mendeleev, its modern formulation. What is its difference from the one given by D.I. Mendeleev? Explain what caused this change in the wording of the law? What is physical meaning Periodic law? Explain the reason for periodic changes in properties chemical elements. How do you understand the phenomenon of periodicity?
The periodic law was formulated by D.I. 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 the simple and complex bodies they form, are periodically dependent on their atomic weight.”
Currently, D. I. 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 they form, 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 does not have an 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 N.D. 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, a weakening of metallic properties and an increase in non-metallic properties are observed. The final element of the period is a noble gas. Each subsequent period begins with an alkali metal, i.e., as the atomic mass of the elements increases, the change chemical properties has a periodic character.
With the development of atomic physics and quantum chemistry, the periodic law received strict theoretical basis. Thanks to the classic works of J. Rydberg (1897), A. Van den Broek (1911), G. Moseley (1913), the physical meaning of the serial (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 increase (N. Bohr, W. Pauli, E. Schrödinger, W. Heisenberg, etc.).
Periodic properties of chemical elements
In principle, the properties of a chemical element combine 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 it forms. 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 exhibit a clear periodic dependence on the atomic numbers of chemical elements. Among these properties, the most important, and of particular importance in explaining or predicting the chemical behavior of elements and the compounds they form, are:
Ionization energy of atoms;
Electron affinity energy of atoms;
Electronegativity;
Atomic (and ionic) radii;
Energy of atomization of simple substances
Oxidation states;
Oxidation 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 similar electronic structures of atoms periodically renewed at increasingly higher energy levels. With their regular change, the physical and chemical properties naturally change.
The physical meaning of the periodic law became clear after the creation of the theory of atomic structure.
So, the physical meaning of the periodic law is that the periodic change in the properties of elements is in full accordance with the similar electronic structures of atoms periodically renewed at ever higher energy levels. With their regular change, the physical and chemical properties of the elements naturally change.
What is the physical meaning of the periodic law.
These conclusions reveal the physical meaning of D.I. Mendeleev’s periodic law, which remained unclear for half a century after the discovery of this law.
It follows that the physical meaning of D.I. Mendeleev’s periodic law consists in the periodic repetition of similar electronic configurations with an increase in the principal quantum number and the unification of elements according to the proximity of their electronic structure.
The theory of atomic structure has shown that the physical meaning of the periodic law is that with a successive increase in nuclear charges, similar valence electronic structures of atoms are periodically repeated.
From all of the above, it is clear that the theory of the structure of the atom revealed the physical meaning of D. I. Mendeleev’s periodic law and even more clearly revealed its significance as the basis for further development chemistry, physics and a number of other sciences.
Replacing the atomic mass with the charge of the nucleus 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 of the periodic function of the dependence of properties on the charge of the nucleus, explain the values of the periods, the number of rare earth elements, etc.
For analogue elements, the same number of electrons is observed in shells of the same name at different meanings 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 renewed similar electron shells of atoms with a consistent increase in the values of the principal quantum number.
For analogue elements, the same number of electrons is observed in the orbitals of the same name at different values of 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 renewed similar electron shells of atoms with a consistent increase in the values of the principal quantum number.
Thus, with a consistent increase in the charges of atomic nuclei, the configuration of the electron shells periodically repeats and, as a consequence, the chemical properties of the elements periodically repeat. This is the physical meaning of the periodic law.
The periodic law of D.I. Mendeleev is the basis of modern chemistry. The study of the structure of atoms reveals the physical meaning of the periodic law and explains the patterns 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 causes of 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 history periodic ecosystem
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4.1. Periodic law D.I. Mendeleev
The periodic law is the greatest achievement of chemical science, the basis of all modern chemistry. With his discovery, chemistry ceased to be a descriptive science; scientific foresight became possible in it.
Periodic law discovered D. I. Mendeleev in 1869, the Scientist formulated this law as follows: “The properties of simple bodies, as well as the forms and properties of compounds of elements, are periodically dependent on the magnitude of the atomic weights of the elements.”
A more detailed study of the structure of matter showed that the periodicity of the properties of elements is determined not by atomic mass, but by the electronic structure of atoms.
The nuclear charge is a characteristic that determines the electronic structure of atoms, and therefore the properties of elements. Therefore, in the modern formulation, the Periodic Law sounds like this: the properties of simple substances, as well as the forms and properties of compounds of elements, are periodically dependent on the atomic number (on the charge value of the nucleus of their atoms).
The expression of the Periodic Law is the periodic table of elements.
4.2. Periodic table of D. I. Mendeleev
The periodic table of elements by D.I. Mendeleev consists of seven periods, which are horizontal sequences of elements arranged in increasing order of the charge of their atomic nucleus. Periods 1, 2, 3, 4, 5, 6 contain 2, 8, 8, 18, 18, 32 elements, respectively. The seventh period is not completed. Periods 1, 2 and 3 are called small, the rest - big.
Each period (except for the first) begins with atoms of alkali metals (Li, Na, K, Rb, Cs, Fr) and ends with a noble gas (Ne, Ar, Kr, Xe, Rn), which is preceded by a typical non-metal. In periods from left to right, metallic ones gradually weaken and non-metallic ones intensify. metallic properties, since as the positive charge of atomic nuclei increases, the number of electrons at the external level increases.
In the first period, besides helium, there is only one element - hydrogen. It is conditionally placed in subgroup IA or VIIA, since it shows similarities with both alkali metals and halogens. The similarity of hydrogen with alkali metals is manifested in the fact that hydrogen, like alkali metals, is a reducing agent and, by donating one electron, forms a singly charged cation. Hydrogen has more in common with halogens: hydrogen, like halogens, is a non-metal, its molecule is diatomic, it can exhibit oxidizing properties, forming salt-like hydrides with active metals, for example, NaH, CaH 2.
In the fourth period, following Ca, there are 10 transition elements (decade Sc - Zn), followed by the remaining 6 main elements of the period (Ga - Kg). The fifth period is constructed similarly. Concept transition element usually used to refer to any element with d or f valence electrons.
The sixth and seventh periods have double insertions of elements. Behind the Ba element there is an inserted decade of d-elements (La - Hg), and after the first transition element La there are 14 f-elements - lanthanides(Se - Lu). After Hg there are the remaining 6 main p-elements of the sixth period (Tl - Rn).
In the seventh (incomplete) period, Ac is followed by 14 f-elements- actinides(Th - Lr). IN lately La and Ac began to be classified as lanthanides and actinides, respectively. Lanthanides and actinides are placed separately at the bottom of the table.
Thus, each element in the periodic table occupies a strictly defined position, which is marked ordinal, or atomic number.
In the periodic table, eight groups are located vertically (I – VIII), which in turn are divided into subgroups - main ones, or subgroups A and side effects, or subgroup B. Subgroup VIIIB is special, it contains triads elements that make up the iron family (Fe, Co, Ni) and platinum metals(Ru, Rh, Pd, Os, Ir, Pt).
The similarity of elements within each subgroup is the most noticeable and important pattern in the periodic table. In the main subgroups, from top to bottom, metallic properties increase and non-metallic properties weaken. In this case, there is an increase in the stability of compounds of elements in the lowest oxidation state for a given subgroup. In side subgroups, on the contrary, from top to bottom, the metallic properties weaken and the stability of compounds with the highest oxidation state increases.
4.3. Periodic table and electronic configurations of atoms
Since when chemical reactions the nuclei of the reacting atoms do not change, then the chemical properties of the atoms depend on the structure of their electronic shells.
The filling of electronic layers and electron shells of atoms occurs in accordance with the Pauli principle and Hund's rule.
Pauli's principle (Pauli's exclusion)
Two electrons in an atom cannot have four identical quantum numbers (on each atomic orbital there can be no more than two electrons).
The Pauli principle determines the maximum number of electrons possessing a given principal quantum number n(i.e. located on this electronic layer): N n = 2n 2. The first electron layer (energy level) can have no more than 2 electrons, the second – 8, the third – 18, etc.
In a hydrogen atom, for example, there is one electron, which is in the first energy level in the 1s state. The spin of this electron can be directed arbitrarily (m s = +1/2 or m s = –1/2). It should be emphasized once again that the first energy level consists of one sublevel - 1s, the second energy level - of two sublevels - 2s and 2p, the third - of three sublevels - 3s, 3p, 3d, etc. The sublevel, in turn, contains orbitals, the number of which is determined by the side quantum number l and equals (2 l + 1). Each orbital is conventionally designated by a square, the electron located on it is designated by an arrow, the direction of which indicates the orientation of the spin of this electron. This means that the state of an electron in a hydrogen atom can be represented as 1s 1 or depicted as a quantum cell, Fig. 4.1:
Rice. 4.1. Symbol electron in a hydrogen atom in the 1s orbital
For both electrons of the helium atom n = 1, l = 0, m l= 0, m s = +1/2 and –1/2. Therefore, the electronic formula of helium is 1s 2. The electron shell of helium is complete and very stable. Helium is a noble gas.
According to the Pauli principle, there cannot be two electrons with parallel spins in one orbital. The third electron in a lithium atom occupies the 2s orbital. The electronic configuration of Li is 1s 2 2s 1, and that of beryllium is 1s 2 2s 2. Since the 2s orbital is filled, the fifth electron of the boron atom occupies the 2p orbital. At n= 2 side (orbital) quantum number l takes values 0 and 1. When l = 0 (2s-state) m l= 0, and at l = 1 (2p – state) m l may be equal to +1; 0; –1. The 2p state corresponds to three energy cells, Fig. 4.2.
Rice. 4.2. Arrangement of electrons of a boron atom in orbitals
For the nitrogen atom (electronic configuration 1s 2 2s 2 2p 3 two electrons on the first level, five on the second), the following two options for the electronic structure are possible, Fig. 4.3:
Rice. 4.3. Possible options arrangement of electrons of the nitrogen atom in orbitals
In the first scheme, Fig. 4.3a, the total spin is equal to 1/2 (+1/2 –1/2 +1/2), in the second (Fig. 4.3b) the total spin is equal to 3/2 (+1/2 + 1/2 +1/2). The location of the spins is determined Hund's rule which reads: filling of energy levels occurs in such a way that the total spin is maximum.
Thus , Of the two given schemes for the structure of the nitrogen atom, the first one corresponds to the stable state (with the lowest energy), where all p-electrons occupy different orbitals. The sublevel orbitals are filled as follows: first, one electron with the same spins, and then a second electron with opposite spins.
Starting with sodium, the third energy level with n = 3 is filled. The distribution of electrons of atoms of elements of the third period in orbitals is shown in Fig. 4.4.
Rice. 4.4. Distribution of electrons in orbitals for atoms of elements of the third period in the ground state
In an atom, each electron occupies a free orbital with the lowest energy corresponding to its strongest connection with the nucleus. In 1961 V.M. Klechkovsky formulated general position, according to which the energy of electron orbitals increases in the order of increasing the sum of the main and secondary quantum numbers ( n + l), and in the case of equality of these sums, the orbital with a lower value of the principal quantum number n has less energy.
The sequence of energy levels in order of increasing energy is approximately as follows:
1s< 2s < 2p < 3s < 3р < 4s ≈ 3d < 4p < 5s ≈ 4d < 5p < 6s ≈ 5d ≈ 4f < 6p.
Let's consider the distribution of electrons in the orbitals of atoms of elements of the fourth period (Fig. 4.5).
Rice. 4.5. Distribution of electrons over the orbitals of atoms of elements of the fourth period in the ground state
After potassium (electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1) and calcium (electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2), the inner 3d shell is filled with electrons (transition elements Sc - Zn) . It should be noted that there are two anomalies: for the Cr and Cu atoms at 4 s-shell contains not two electrons, but one, i.e. the so-called “failure” of the outer 4s electron to the previous 3d shell occurs. The electronic structure of the chromium atom can be represented as follows (Fig. 4.6).
Rice. 4.6. Distribution of electrons over orbitals for the chromium atom
The physical reason for the “violation” of the filling order is associated with the different penetrating ability of electron orbitals to the nucleus, the special stability of the electronic configurations d 5 and d 10, f 7 and f 14, corresponding to the filling of electronic orbitals with one or two electrons, as well as the screening effect of the internal electronic charge layers kernels.
The electronic configurations of Mn, Fe, Co, Ni, Cu and Zn atoms are reflected by the following formulas:
25 Mn 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 2,
26 Fe 1s 2 2s 2 2p 6 3s 2 3p 6 3d 6 4s 2,
27 Co 1s 2 2s 2 2p 6 3s 2 3p 6 3d 7 4s 2,
28 Ni 1s 2 2s 2 2p 6 3s 2 3p 6 3d 8 4s 2,
29 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 ,
30 Zn 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 .
After zinc, starting from the 31st element - gallium up to the 36th element - krypton, the filling of the fourth layer (4p - shell) continues. The electronic configurations of these elements are as follows:
31 Ga 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 1 ,
32 Ge 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 2 ,
33 As 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 3 ,
34 Se 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 4,
35 Br 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 5,
36 Kr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 .
It should be noted that if the Pauli exclusion is not violated, in excited states electrons can be located in other atomic orbitals.
4.4. Types of chemical elements
All elements of the periodic table are divided into four types:
1. In atoms s-elements the s-shells of the outer layer (n) are filled. The s elements include hydrogen, helium and the first two elements of each period.
2. At atoms p-elements electrons fill the p-shells of the outer level (np). The p-elements include the last 6 elements of each period (except the first).
3. U d-elements is filled with electrons d–shell of the second outer level (n–1) d. These are elements of plug-in decades of large periods located between the s- and p- elements.
4. U f-elements the f sublevel of the third outside level (n–2) f is filled with electrons. The family of f-elements includes lanthanides and actinides.
From considering the electronic structure of unexcited atoms depending on the atomic number of the element, it follows:
The number of energy levels (electronic layers) of an atom of any element is equal to the number of the period in which the element is located. This means that s-elements are found in all periods, p-elements in the second and subsequent periods, d-elements in the fourth and subsequent periods, and f-elements in the sixth and seventh periods.
The period number coincides with the principal quantum number of the outer electrons of the atom.
s- and p-elements form the main subgroups, d-elements form secondary subgroups, f-elements form the families of lanthanides and actinides. Thus, the subgroup includes elements whose atoms usually have a similar structure not only of the outer, but also of the pre-outer layer (with the exception of elements in which there is a “failure” of the electron).
The group number usually indicates the number of electrons that can participate in the formation of chemical bonds. This is the physical meaning of the group number. Elements of side subgroups have valence electrons not only in their outer shells, but also in their penultimate shells. This is the main difference in the properties of the elements of the main and secondary subgroups.
Elements with valence d- or f-electrons are called transition elements.
The group number, as a rule, is equal to the highest positive oxidation state of the elements that they exhibit in compounds. The exception is fluorine - its oxidation state is –1; from elements VIII group only for Os, Ru and Xe the oxidation state +8 is known.
4.5. Periodicity of properties of atoms of elements
Such characteristics of atoms as their radius, ionization energy, electron affinity, electronegativity, and oxidation state are associated with the electronic structure of the atom.
There are radii of metal atoms and covalent radii of non-metal atoms. The radii of metal atoms are calculated based on interatomic distances, which are well known for most metals based on experimental data. In this case, the radius of a metal atom is equal to half the distance between the centers of two neighboring atoms. The covalent radii of nonmetals in molecules and crystals of simple substances are calculated in a similar way. The larger the atomic radius, the easier it is for outer electrons to break away from the nucleus (and vice versa). Unlike atomic radii, ion radii are arbitrary values.
From left to right in periods, the value of the atomic radii of metals decreases, and the atomic radii of non-metals changes in a complex way, since it depends on the nature of the chemical bond. In the second period, for example, the radii of the atoms first decrease and then increase, especially sharply when moving to a noble gas atom.
In the main subgroups, the radii of atoms increase from top to bottom, as the number of electronic layers increases.
The radius of a cation is less than the radius of its corresponding atom, and as the positive charge of the cation increases, its radius decreases. On the contrary, the radius of an anion is always greater than the radius of its corresponding atom. Particles (atoms and ions) that have the same number of electrons are called isoelectronic. In the series of isoelectronic ions, the radius decreases as the negative radius of the ion decreases and the positive radius increases. Such a decrease occurs, for example, in the series: O 2–, F–, Na +, Mg 2+, Al 3+.
Ionization energy– the energy required to remove an electron from an atom in the ground state. It is usually expressed in electron volts (1 eV = 96.485 kJ/mol). In a period, from left to right, the ionization energy increases with increasing nuclear charge. In the main subgroups, from top to bottom, it decreases, since the distance of the electron to the nucleus increases and the screening effect of the internal electronic layers increases.
Table 4.1 shows the values of ionization energies (energies for the removal of the first, second, etc. electrons) for some atoms.
In the second period, during the transition from Li to Ne, the energy of removal of the first electron increases (see Table 4.1). However, as can be seen from the table, the ionization energy increases unevenly: for boron and oxygen, which follow beryllium and nitrogen, respectively, a slight decrease is observed, which is due to the peculiarities of the electronic structure of atoms.
The outer s-shell of beryllium is completely filled, so the electron next to it, boron, enters the p-orbital. This p-electron is less tightly bound to the nucleus than the s-electron, so the removal of p-electrons requires less energy.
Table 4.1.
Ionization energies I atoms of some elements
Each p-orbital of the nitrogen atom has one electron. In the oxygen atom, an electron enters the p-orbital, which is already occupied by one electron. Two electrons in the same orbital repel strongly, so it is easier to remove an electron from an oxygen atom than from a nitrogen atom.
Alkali metals have the lowest ionization energy, therefore they have pronounced metallic properties; the highest ionization energy is for inert gases.
Electron affinity– energy released when an electron attaches to a neutral atom. Electron affinity, like ionization energy, is usually expressed in electron volts. The highest electron affinity is for halogens, the lowest for alkali metals. Table 4.2 shows the electron affinities for atoms of some elements.
Table 4.2.
Electron affinities of atoms of some elements
Electronegativity- the ability of an atom in a molecule or ion to attract valence electrons from other atoms. Electronegativity (EO) as a quantitative measure is an approximate value. About 20 electronegativity scales have been proposed, the most widely recognized of which is the scale developed by L. Pauling. In Fig. 4.7 shows the values of EO according to Pauling.
Rice. 4.7. Electronegativity of elements (according to Pauling)
Fluorine is the most electronegative of all elements on the Pauling scale. Its EO is taken to be 4. The least electronegative is cesium. Hydrogen occupies an intermediate position, since when interacting with some elements it gives up an electron, and when interacting with others it gains.
4.6. Acid-base properties of compounds; Kossel circuit
To explain the nature of changes in the acid-base properties of compounds of elements, Kossel (Germany) proposed using simple diagram, based on the assumption that there is a purely ionic bond in the molecules and a Coulomb interaction takes place between the ions. The Kossel scheme describes the acid-base properties of compounds containing E-N connections and E-O-N, depending on the charge of the nucleus and the radius of the element forming them.
The Kossel diagram for two metal hydroxides, such as LiOH and KOH, is shown in Fig. 4.8.
Rice. 4.8. Kossel diagram for LiOH and KOH
As can be seen from the presented diagram, the radius of the Li + ion is smaller than the radius of the K + ion and the OH - group is bound more tightly to the lithium cation than to the potassium cation. As a result, KOH will be easier to dissociate in solution and the basic properties of potassium hydroxide will be more pronounced.
The Kossel diagram for two bases CuOH and Cu(OH) 2 can be analyzed in a similar way. Since the radius of the Cu 2+ ion is smaller and the charge is greater than that of the Cu + ion, the OH - group will be held more firmly by the Cu 2+ ion. As a result, the base Cu(OH) 2 will be weaker than CuOH.
Thus, the strength of bases increases as the radius of the cation increases and its positive charge decreases.
In the main subgroups, from top to bottom, the strength of the bases increases as the radii of the element ions increase in this direction. In periods from left to right, the radii of the element ions decrease and their positive charge increases, so the strength of the bases decreases in this direction.
The Kossel diagram for two oxygen-free acids, for example, HCl and HI, is shown in Fig. 4.9
Rice. 4.9. Kossel diagram for HCl and HI
Since the radius of the chloride ion is smaller than that of the iodide ion, the H+ ion is more strongly bound to the anion in the hydrochloric acid molecule, which will be weaker than hydroiodic acid. Thus, the strength of anoxic acids increases with increasing radius of the negative ion.
The strength of oxygen-containing acids changes in the opposite way. It increases as the radius of the ion decreases and its positive charge increases. In Fig. Figure 4.10 shows the Kossel diagram for two acids HClO and HClO 4.
Rice. 4.10. Kossel diagram for HClO and HClO 4
The C1 7+ ion is firmly bonded to the oxygen ion, so the proton will be more easily split off in the HC1O 4 molecule. At the same time, the bond between the C1+ ion and the O2- ion is less strong, and in the HC1O molecule the proton will be more strongly retained by the O2- anion. As a result, HClO 4 will be more strong acid than HClO.
The advantage of Kossel's scheme is that, using a simple model, it allows one to explain the nature of changes in the acid-base properties of compounds in a series of similar substances. However, this scheme is purely qualitative. It only allows you to compare the properties of compounds and does not make it possible to determine the acid-base properties of an arbitrarily selected single compound. The disadvantage of this model is that it is based only on electrostatic concepts, while in nature there is no pure (one hundred percent) ionic bond.
4.7. Redox properties of elements and their compounds
A change in the redox properties of simple substances can be easily established by considering the nature of the change in the electronegativity of the corresponding elements. In the main subgroups, from top to bottom, electronegativity decreases, which leads to a decrease in oxidative properties and an increase in reducing properties in this direction. In periods from left to right, electronegativity increases. As a result, in this direction, the reducing properties of simple substances decrease, and the oxidizing properties increase. Thus, strong reducing agents are located in the lower left corner of the periodic table of elements (potassium, rubidium, cesium, barium), while strong oxidizing agents are located in its upper right corner (oxygen, fluorine, chlorine).
The redox properties of compounds of elements depend on their nature, the degree of oxidation of the elements, the position of the elements in the periodic table and a number of other factors.
In the main subgroups, from top to bottom, the oxidizing properties of oxygen-containing acids, in which the atoms of the central element have the same oxidation state, decrease. Strong oxidizing agents are nitric and concentrated sulfuric acids. The greater the positive oxidation state of the element in the compound, the more pronounced its oxidizing properties are. Potassium permanganate and potassium dichromate exhibit strong oxidizing properties.
In the main subgroups, the reducing properties of simple anions increase from top to bottom. Strong reducing agents are HI, H 2 S, iodides and sulfides.
The periodic law of chemical elements is a fundamental law of nature that establishes the periodicity of changes in the properties of chemical elements as the charges of the nuclei of their atoms increase. The date of discovery of the law is considered to be March 1 (February 17, old style) 1869, when D. I. Mendeleev completed the development of the “Experience of a system of elements based on their atomic weight and chemical similarity.” The scientist first used the term “periodic law” (“law of periodicity”) at the end of 1870. According to Mendeleev, “three types of data” contributed to the discovery of the periodic law. Firstly, availability is sufficient large number known elements(63); secondly, satisfactory knowledge of the properties of most of them; thirdly, the fact that the atomic weights of many elements were determined with good accuracy, thanks to which chemical elements could be arranged in a natural series according to the increase in their atomic weights. Mendeleev considered the decisive condition for the discovery of the law to be the comparison of all elements according to their atomic weights (previously only chemically similar elements were compared).
The classic formulation of the periodic law, given by Mendeleev in July 1871, stated: “The properties of the elements, and therefore the properties of the simple and complex bodies they form, are periodically dependent on their atomic weight.” This formulation remained in force for more than 40 years, but the periodic law remained only a statement of facts and had no physical basis. It became possible only in the mid-1910s, when the nuclear planetary model of the atom was developed (see Atom) and it was established that the serial number of an element in the periodic table is numerically equal to the charge of the nucleus of its atom. As a result, the physical formulation of the periodic law became possible: “Properties of elements and the simple and complex substances are periodically dependent on the nuclear charges (Z) of their atoms.” It is still widely used today. The essence of the periodic law can be expressed in other words: “The configurations of the outer electron shells of atoms are periodically repeated as Z increases”; This is a kind of “electronic” formulation of the law.
An essential feature of the periodic law is that, unlike some other fundamental laws of nature (for example, the law of universal gravitation or the law of equivalence of mass and energy), it does not have a quantitative expression, that is, it cannot be written in the form of any or a mathematical formula or equation. Meanwhile, Mendeleev himself and other scientists tried to look for a mathematical expression of the law. In the form of formulas and equations, various patterns of constructing electronic configurations of atoms can be quantitatively expressed depending on the values of the principal and orbital quantum numbers. As for the periodic law, it has a clear graphical reflection in the form of a periodic system of chemical elements, represented mainly various types tables (see insert).
The periodic law is a universal law for the entire Universe, manifesting itself wherever material structures of the atomic type exist. However, it is not only the configurations of atoms that periodically change as Z increases. It turned out that the structure and properties of atomic nuclei also change periodically, although the very nature of the periodic change here is much more complicated than in the case of atoms: in the nuclei there is a regular formation of proton and neutron shells. Nuclei in which these shells are filled (they contain 2, 8, 20, 50, 82, 126 protons or neutrons) are called “magic” and are considered as a kind of boundaries of the periods of the periodic system of atomic nuclei.
DI. Mendeleev formulated the Periodic Law in 1869, which was based on one of the main characteristics atom - atomic mass. The subsequent development of the Periodic Law, namely, the acquisition of a large amount of experimental data, somewhat changed the original formulation of the law, but these changes do not contradict the main meaning laid down by D.I. Mendeleev. These changes only gave the law and the Periodic Table scientific validity and confirmation of correctness.
Modern formulation of the Periodic Law by D.I. Mendeleev is as follows: the properties of chemical elements, as well as the properties and forms of compounds of elements, are periodically dependent on the magnitude of the charge of the nuclei of their atoms.
Structure of the Periodic Table of Chemical Elements D.I. Mendeleev
By present opinion it is known large number interpretations of the Periodic Table, but the most popular is with short (small) and long (large) periods. Horizontal rows are called periods (they contain elements with sequential filling of the same energy level), and vertical columns are called groups (they contain elements that have the same number of valence electrons - chemical analogues). Also, all elements can be divided into blocks according to the type of external (valence) orbital: s-, p-, d-, f-elements.
There are a total of 7 periods in the system (table), and the period number (indicated Arabic numeral) is equal to the number of electronic layers in an element’s atom, the number of the external (valence) energy level, and the value of the principal quantum number for the highest energy level. Each period (except the first) begins with an s-element - an active alkali metal and ends with an inert gas, preceded by a p-element - an active non-metal (halogen). If you move through the period from left to right, then with an increase in the charge of the nuclei of atoms of chemical elements of small periods, the number of electrons at the external energy level will increase, as a result of which the properties of the elements change - from typically metallic (since at the beginning of the period there is an active alkali metal), through amphoteric (the element exhibits the properties of both metals and non-metals) to non-metallic (the active non-metal is halogen at the end of the period), i.e. metallic properties gradually weaken and non-metallic properties increase.
In large periods, as the charge of nuclei increases, the filling of electrons is more difficult, which explains a more complex change in the properties of elements compared to elements of small periods. Thus, in even rows of long periods, with increasing charge of the nucleus, the number of electrons in the outer energy level remains constant and equal to 2 or 1. Therefore, while the level next to the outer (second from the outside) is filled with electrons, the properties of the elements in the even rows change slowly. When moving to odd series, with increasing nuclear charge, the number of electrons in the external energy level increases (from 1 to 8), the properties of the elements change in the same way as in small periods.
Vertical columns in the Periodic Table are groups of elements with similar electronic structures and which are chemical analogues. Groups are designated by Roman numerals from I to VIII. There are main (A) and secondary (B) subgroups, the first of which contain s- and p-elements, the second - d-elements.
The number A of the subgroup shows the number of electrons in the outer energy level (the number of valence electrons). For B-subgroup elements, there is no direct connection between the group number and the number of electrons in the outer energy level. In A-subgroups, the metallic properties of elements increase, and non-metallic properties decrease with increasing charge of the nucleus of the element’s atom.
There is a relationship between the position of elements in the Periodic Table and the structure of their atoms:
- atoms of all elements of the same period have an equal number of energy levels, partially or completely filled with electrons;
- atoms of all elements of the A subgroups have an equal number of electrons at the outer energy level.
Periodic properties of elements
The similarity of the physicochemical and chemical properties of atoms is due to the similarity of their electronic configurations, and main role plays the distribution of electrons in the outer atomic orbital. This manifests itself in the periodic appearance, as the charge of the atomic nucleus increases, of elements with similar properties. Such properties are called periodic, among which the most important are:
1. Number of electrons in the outer electron shell ( population – w). In short periods with increasing nuclear charge w the outer electron shell monotonically increases from 1 to 2 (1st period), from 1 to 8 (2nd and 3rd periods). In large periods during the first 12 elements w does not exceed 2, and then up to 8.
2. Atomic and ionic radii(r), defined as the average radii of an atom or ion, found from experimental data on interatomic distances in different compounds. By period, the atomic radius decreases (gradually adding electrons are described by orbitals with almost equal characteristics; by group, the atomic radius increases as the number of electron layers increases (Fig. 1).
Rice. 1. Periodic change in atomic radius
The same patterns are observed for the ionic radius. It should be noted that the ionic radius of the cation (positively charged ion) is greater than the atomic radius, which in turn is greater than the ionic radius of the anion (negatively charged ion).
3. Ionization energy(E and) is the amount of energy required to remove an electron from an atom, i.e. the energy required to transform a neutral atom into a positively charged ion (cation).
E 0 - → E + + E and
E and is measured in electronvolts (eV) per atom. Within the group of the Periodic Table, the values of ionization energy of atoms decrease with increasing charges of the atomic nuclei of elements. All electrons can be sequentially removed from atoms of chemical elements by reporting discrete values of E and. Moreover, E and 1< Е и 2 < Е и 3 <….Энергии ионизации отражают дискретность структуры электронных слоев и оболочек атомов химических элементов.
4. Electron affinity(E e) – the amount of energy released when an additional electron is added to an atom, i.e. process energy
E 0 + → E —
E e is also expressed in eV and, like E, it depends on the radius of the atom, therefore the nature of the change in E e across periods and groups of the Periodic System is close to the nature of the change in the atomic radius. Group VII p-elements have the highest electron affinity.
5. Regenerative activity(VA) – the ability of an atom to give an electron to another atom. Quantitative measure – E and. If E increases, then BA decreases and vice versa.
6. Oxidative activity(OA) – the ability of an atom to attach an electron from another atom. Quantitative measure E e. If E e increases, then OA also increases and vice versa.
7. Shielding effect– reducing the impact of the positive charge of the nucleus on a given electron due to the presence of other electrons between it and the nucleus. Shielding increases with the number of electron layers in an atom and reduces the attraction of outer electrons to the nucleus. The opposite of shielding penetration effect, due to the fact that the electron can be located at any point in atomic space. The penetration effect increases the strength of the bond between the electron and the nucleus.
8. Oxidation state (oxidation number)– the imaginary charge of an atom of an element in a compound, which is determined from the assumption of the ionic structure of the substance. The group number of the Periodic Table indicates the highest positive oxidation state that elements of a given group can have in their compounds. Exceptions are metals of the copper subgroup, oxygen, fluorine, bromine, metals of the iron family and other elements of group VIII. As the nuclear charge increases in a period, the maximum positive oxidation state increases.
9. Electronegativity, compositions of higher hydrogen and oxygen compounds, thermodynamic, electrolytic properties, etc.
Examples of problem solving
EXAMPLE 1
| Exercise | Characterize the element (Z=23) and the properties of its compounds (oxides and hydroxides) using the electronic formula: family, period, group, number of valence electrons, electron graphic formula for valence electrons in the ground and excited states, main oxidation states (maximum and minimum ), formulas of oxides and hydroxides. |
| Solution | 23 V 1s 2 2s 2 2p 6 3s 3 3p 6 3d 3 4s 2 d-element, metal, is in the ;-th period, in the V group, in the subgroup. Valence electrons 3d 3 4s 2. Oxides VO, V 2 O 3, VO 2, V 2 O 5. Hydroxides V(OH)2, V(OH)3, VO(OH)2, HVO3. Ground state Excited state The minimum oxidation state is “+2”, the maximum is “+5”. |
2.3. Periodic law of D.I.Mendeleev.
The law was discovered and formulated by D.I. 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 achievements in physics, mainly the development of the theory of atomic structure, made it possible to reveal the physical essence of the periodic law: the periodicity in changes in the properties of chemical elements is due to a periodic change in the nature of the filling of the outer electron layer with electrons as the number of electrons, determined by the charge of the nucleus, increases. The charge is equal to the atomic number of the element in the periodic table. The modern formulation of the periodic law: “The properties of elements and the simple and complex substances they form are periodically dependent on the charge of the nuclei of atoms.” 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 precisely formulate this law and present its contents in the form of a table, which became classic, but also comprehensively substantiated it, showed its enormous scientific significance, as a guiding classification principle and as a powerful tool for scientific research.
Physical meaning of the periodic law. It was opened only after it was discovered that the charge of the nucleus of an atom increases when moving from one chemical element to a neighboring one (in the periodic table) by a unit of elementary charge. Numerically, the charge of the nucleus is equal to the atomic number (atomic number Z) of the corresponding element in the periodic table, 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 changes with increasing nuclear charge, and, therefore, the periodic law is based on the idea of a change in the charge of the nucleus of atoms, and not the atomic mass of the elements. A clear illustration of the periodic law is the curves of periodic changes in certain physical quantities (ionization potentials, atomic radii, atomic volumes) depending on Z. There is no general mathematical expression for the periodic law. The periodic law has enormous natural scientific and philosophical significance. It made it possible to consider all elements in their mutual connection 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 clear manifestation of the general laws of dialectics, in particular the law of the transition of quantity into 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 atomic structure (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 element number in the periodic table;
5. Development of the theory of the periodic system based on ideas about the structure of the electronic 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 chemical elements that have not yet been discovered. Under aluminum Al, Mendeleev left a place for its analogue “eka-aluminium”, under boron B - for “eca-boron”, and under silicon Si - for “eca-silicon”. This is what Mendeleev called the yet undiscovered chemical elements. He even gave them the symbols El, Eb and Es.
Regarding the element “exasilicon,” 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 analogues, namely, to the III row. This will be the metal immediately following silicon, and therefore let us call it ekasilicium." 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 “dwitellurium” - polonium (discovered in 1898), “ecaiod” - astatine (discovered in 1942-1943), “dimanganese” - technetium (discovered in 1937) , "ecacesia" - France (opened in 1939)
In 1875, the French chemist Paul-Emile Lecoq de Boisbaudran discovered the “eka-aluminum” predicted by Mendeleev in the mineral wurtzite - zinc sulfide ZnS - and named it gallium Ga (the Latin name for France is “Gallia”) in honor of his homeland.
Mendeleev accurately predicted the properties of eka-aluminum: its atomic mass, the density of the metal, the formula of El 2 O 3 oxide, ElCl 3 chloride, El 2 (SO 4) 3 sulfate. 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 fusible metal, and indeed, the melting point of gallium turned out to be equal to 29.8 o C. In terms of fusibility, gallium is second only to mercury Hg and cesium Cs.
The average content of gallium in the earth's crust is relatively high, 1.5-10-30% by mass, which is equal to the content of lead and molybdenum. Gallium is a typical trace element. The only Gallium mineral is galdite CuGaS2, which is very rare. Gallium is stable in air at ordinary temperatures. Above 260°C, slow oxidation is observed in dry oxygen (the oxide film protects the metal). Gallium dissolves slowly in sulfuric and hydrochloric acids, quickly in hydrofluoric acid, and is stable in the cold in nitric acid. Gallium dissolves slowly 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 the large range of the liquid state (2200° C) and low vapor pressure at temperatures up to 1100-1200° C. Geochemistry Gallium is closely related to the geochemistry of aluminum, which is due to the similarity of their physicochemical properties. The main part 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 hard coals (together with germanium), and also in some iron ores. Gallium does not yet have widespread industrial use. The potential scale of by-product 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, which have semiconductor properties. They can be used in high-temperature rectifiers and transistors, solar batteries and other devices where the photoelectric effect in the blocking layer can be used, as well as in infrared radiation receivers. Gallium can be used to make optical mirrors that are highly reflective. An alloy of aluminum with gallium has been proposed instead of mercury as the cathode of ultraviolet radiation lamps used in medicine. It is proposed to use liquid gallium and its alloys for the manufacture of high-temperature thermometers (600-1300 ° C) and pressure gauges. Of interest is the use of Gallium and its alloys as a liquid coolant in power nuclear reactors (this is hampered by the active interaction of Gallium at operating temperatures with structural materials; the eutectic Ga-Zn-Sn alloy has a less corrosive effect than pure Gallium).
In 1879, Swedish chemist Lars Nilsson discovered scandium, predicted by Mendeleev as ecaboron Eb. Nilsson wrote: “There remains no doubt that ecaboron 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 in honor of Nilsson'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 mass. Scandium content in rocks varies: in ultrabasic rocks 5-10-4, in basic rocks 2.4-10-3, in intermediate rocks 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 supergene (surface) processes. Two of Scandium's own minerals are known - tortveitite and sterrettite; they are extremely rare. Scandium is a soft metal, in its pure state it can be easily processed - forged, rolled, stamped. The scope of use of scandium is very limited. Scandium oxide is used to make ferrites for memory elements of high-speed computers. 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 find application in the manufacture of phosphors, oxide cathodes, in glass and ceramic production, in the chemical industry (as catalysts) and in others areas. In 1886, a professor at the Mining Academy in Freiburg, the German chemist Clemens Winkler, while analyzing the rare mineral argyrodite with the composition Ag 8 GeS 6, discovered another element predicted by Mendeleev. Winkler named the element he discovered germanium Ge in honor of his homeland, but for some reason this caused sharp 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 “ekasilicium” and the symbol Es. Discouraged, Winkler turned to Dmitry Ivanovich himself for advice. He explained that it was the discoverer of the new element who should give it a name. The total content of germanium in the earth's crust is 7.10-4% by mass, i.e. more than, for example, antimony, silver, bismuth. However, germanium's own minerals are extremely rare. Almost all of them are sulfosalts: germanite Cu2 (Cu, Fe, Ge, Zn)2 (S, As)4, argyrodite Ag8GeS6, confieldite Ag8(Sn, Ce) S6, etc. The bulk of germanium is scattered in large quantities in the earth’s crust 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 coal and oil deposits. 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 dosimetric instruments and devices that measure the strength of constant and alternating magnetic fields. An important area of application for germanium is infrared technology, in particular the production of infrared radiation detectors operating in the region of 8-14 microns. Many alloys containing germanium, GeO2-based glasses, and other germanium compounds are promising for practical use.
Mendeleev could not predict the existence of a group of noble gases, and at first they did not find a place in the Periodic Table.
The discovery of argon Ar by 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 an allotropic modification of nitrogen and only in 1900, under the pressure of immutable facts, agreed with the presence of a “zero” group of chemical elements in the Periodic Table, which was occupied by other noble gases discovered after argon. Now this group is known as VIIIA.
In 1905, Mendeleev wrote: “Apparently, the future does not threaten the periodic law with destruction, but only promises superstructures and development, although as a Russian they wanted to erase me, especially the Germans.”
The discovery of the Periodic Law accelerated the development of chemistry and the discovery of new chemical elements.
The lyceum exam, at which old Derzhavin blessed young Pushkin. The role of the meter happened to be played by Academician Yu.F. Fritzsche, a well-known specialist in organic chemistry. Candidate's thesis D.I. Mendeleev graduated from the Main Pedagogical Institute in 1855. His thesis "Isomorphism in connection with other relationships of crystalline form to 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. Botkina, I.M. Sechenova, 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...