As a result of the successful development of the material in this chapter, the student should:
know
- modern formulation of the periodic law;
- connection between the structure of the periodic system and the energy sequence of sublevels in multielectron atoms;
- definitions of the concepts "period", "group", "5-elements", "p-elements", "d- elements”, “/-elements”, “ionization energy”, “electron affinity”, “electronegativity”, “van der Waals radius”, “clarke”;
- basic law of geochemistry;
be able to
Describe the structure of the periodic system in accordance with the rules of Klechkovsky;
own
Ideas about the periodic nature of the change in the properties of atoms and the chemical properties of elements, about the features of the long-period version of the periodic system; about the relationship of the abundance of chemical elements with their position in the periodic system, about macro- and microelements in the lithosphere and living matter.
Modern formulation of the periodic law
Periodic law - the most general law of chemistry - was discovered by Dmitry Ivanovich Mendeleev in 1869. At that time, the structure of the atom was not yet known. D. I. Mendeleev made his discovery based on the regular change in the properties of elements with an increase in atomic masses.
After the discovery of the structure of atoms, it became clear that their properties are determined by the structure of the electron shells, which depends on the total number of electrons in the atom. The number of electrons in an atom is equal to the charge of its nucleus. Therefore, the modern formulation of the periodic law is as follows.
The properties of chemical elements and the simple and complex substances they form are in a periodic dependence on the charge of the nucleus of their atoms.
The significance of the periodic law lies in the fact that it is the main tool for systematizing and classifying chemical information, a very important means of interpreting chemical information, a powerful tool for predicting the properties of chemical compounds, and a means of directed search for compounds with predetermined properties.
The periodic law does not have a mathematical expression in the form of equations, it is reflected in a table called periodic system of chemical elements. There are many variants of the tables of the periodic table. The most widely used are the long-period and short-period versions, placed on the first and second color inserts of the book. The main structural unit of the periodic system is the period.
Period with number p called a sequence of chemical elements arranged in ascending order of the charge of the nucleus of an atom, which begins with ^-elements and ends with ^-elements.
In this definition P - period number equal to the main quantum number for the upper energy level in the atoms of all elements of this period. in atoms s-elements 5-sublevels are completed, in atoms p-elements - respectively p-sublevels. The exception to the above definition is the first period, in which there are no p-elements, since at the first energy level (n = 1) there is only 15-level. The periodic table also contains d-elements, whose ^-sublevels are completed, and /-elements, whose /-sublevels are completed.
Alchemists also tried to find a law of nature, on the basis of which it would be possible to systematize the chemical elements. But they lacked reliable and detailed information about the elements. By the middle of the XIX century. knowledge about chemical elements became sufficient, and the number of elements increased so much that a natural need arose in science to classify them. The first attempts to classify elements into metals and non-metals proved to be untenable. The predecessors of D.I. Mendeleev (I.V. Debereiner, J.A. Newlands, L.Yu. Meyer) did a lot to prepare the discovery of the periodic law, but could not comprehend the truth. Dmitry Ivanovich established a connection between the mass of elements and their properties.
Dmitry Ivanovich was born in Tobolsk. He was the seventeenth child in the family. After graduating from a gymnasium in his native city, Dmitry Ivanovich entered the Main Pedagogical Institute in St. Petersburg, after graduating from which he went on a scientific trip abroad with a gold medal for two years. After returning, he was invited to St. Petersburg University. Starting to read lectures on chemistry, Mendeleev did not find anything that could be recommended to students as a teaching aid. And he decided to write a new book - "Fundamentals of Chemistry".
The discovery of the periodic law was preceded by 15 years of hard work. On March 1, 1869, Dmitry Ivanovich planned to leave St. Petersburg for the province on business.
The periodic law was discovered on the basis of the characteristics of the atom - the relative atomic mass .
Mendeleev arranged the chemical elements in ascending order of their atomic masses and noticed that the properties of the elements repeat after a certain interval - a period, Dmitry Ivanovich placed the periods one under the other., so that similar elements were located one under the other - on the same vertical, so the periodic system was built elements.
March 1, 1869 The formulation of the periodic law by D.I. Mendeleev.
The properties of simple substances, as well as the forms and properties of compounds of elements, are in a periodic dependence on the magnitude of the atomic weights of the elements.
Unfortunately, at first there were very few supporters of the periodic law, even among Russian scientists. There are many opponents, especially in Germany and England.
The discovery of the periodic law is a brilliant example of scientific foresight: in 1870, Dmitry Ivanovich predicted the existence of three then unknown elements, which he called ekasilicium, ekaaluminum and ekabor. He was also able to correctly predict the most important properties of the new elements. And after 5 years, in 1875, the French scientist P.E. Lecoq de Boisbaudran, who knew nothing about the work of Dmitry Ivanovich, discovered a new metal, calling it gallium. In a number of properties and the method of discovery, gallium coincided with ekaaluminum predicted by Mendeleev. But his weight was less than predicted. Despite this, Dmitry Ivanovich sent a letter to France, insisting on his prediction.
The scientific world was stunned that Mendeleev's prediction of properties ekaaluminum
turned out to be so accurate. From this moment, the periodic law begins to assert itself in chemistry.
In 1879, L. Nilson in Sweden discovered scandium, which embodied the predicted by Dmitry Ivanovich ekabor
.
In 1886, K. Winkler discovered germanium in Germany, which turned out to be exasilicon
.
But the genius of Dmitry Ivanovich Mendeleev and his discoveries are not only these predictions!
In four places of the periodic system, D. I. Mendeleev arranged the elements out of order of increasing atomic masses:
As early as the end of the 19th century, D.I. Mendeleev wrote that, apparently, the atom consists of other smaller particles. After his death in 1907, it was proved that the atom consists of elementary particles. The theory of the structure of the atom confirmed the correctness of Mendeleev, the permutations of these elements not in accordance with the growth of atomic masses are fully justified.
The modern formulation of the periodic law.
The properties of chemical elements and their compounds are in a periodic dependence on the magnitude of the charge of the nuclei of their atoms, which is expressed in the periodic repetition of the structure of the outer valence electron shell.
And now, more than 130 years after the discovery of the periodic law, we can return to the words of Dmitry Ivanovich, taken as the motto of our lesson: “The future does not threaten the periodic law with destruction, but only a superstructure and development are promised.” How many chemical elements have been discovered so far? And this is far from the limit.
The graphic representation of the periodic law is the periodic system of chemical elements. This is a brief synopsis of the entire chemistry of the elements and their compounds.
Changes in properties in the periodic system with an increase in the value of atomic weights in the period (from left to right):
1. Metallic properties decrease
2. Non-metallic properties increase
3. The properties of higher oxides and hydroxides change from basic through amphoteric to acidic.
4. The valence of elements in the formulas of higher oxides increases from IbeforeVII, and in the formulas of volatile hydrogen compounds decreases from IV beforeI.
Basic principles of construction of the periodic system.|
Comparison sign |
D.I. Mendeleev |
|
1. How is the sequence of elements by numbers established? (What is the basis of PS?) |
The elements are listed in order of increasing relative atomic masses. However, there are exceptions. Ar - K, Co - Ni, Te - I, Th - Pa |
|
2. The principle of combining elements into groups. |
Quality mark. The similarity of the properties of simple substances and the same type of complex. |
|
3. The principle of combining elements into periods. |
First option Periodic table of elements was published by Dmitri Ivanovich Mendeleev in 1869 and was called "The Experience of a System of Elements".
DI. Mendeleev arranged the 63 elements known at that time in ascending order of their atomic masses and obtained a natural series of chemical elements, in which he discovered a periodic recurrence of chemical properties. This series of chemical elements is now known as the Periodic Law (D.I. Mendeleev's formulation):
The properties of simple bodies, as well as the forms and properties of compounds of elements, are in a periodic dependence on the magnitude of the atomic weights of the elements.
The current wording of the law reads as follows:
The properties of chemical elements, simple substances, as well as the composition and properties of compounds are in a periodic dependence on the values of the charges of the nuclei of atoms.
Graphic image periodic law is the periodic table.
The cell of each element indicates its most important characteristics.
Periodic table contains groups and periods.
Group- a column of the periodic system, in which chemical elements are located that have chemical similarity due to identical electronic configurations of the valence layer.
Periodic system of D.I. Mendeleev contains eight groups of elements. Each group consists of two subgroups: main (a) and secondary (b). The main subgroup contains s- and p- elements, in the side - d- elements.
Group names:
I-a Alkali metals.
II-a Alkaline earth metals.
V-a Pnictogens.
VI-a Chalcogens.
VII-a Halogens.
VIII-a Noble (inert) gases.
Period is a sequence of elements written as a string, arranged in order of increasing charges of their nuclei. The period number corresponds to the number of electronic levels in the atom.
The period starts with an alkali metal (or hydrogen) and ends with a noble gas.
|
Parameter |
Down the group |
By period to the right |
|
Core charge |
is increasing |
is increasing |
|
Number of valence electrons |
Does not change |
is increasing |
|
Number of energy levels |
is increasing |
Does not change |
|
Atom radius |
is increasing |
Decreases |
|
Electronegativity |
Decreases |
is increasing |
|
Metal properties |
Are increasing |
Decrease |
|
Oxidation state in higher oxide |
Does not change |
is increasing |
|
The degree of oxidation in hydrogen compounds (for elements of groups IV-VII) |
Does not change |
is increasing |
Modern periodic table of chemical elements of Mendeleev.
SESSION 5 10th grade(first year of study)
Periodic law and the system of chemical elements d.I. Mendeleev Plan
1. The history of the discovery of the periodic law and the system of chemical elements by D.I. Mendeleev.
2. Periodic law in the formulation of DIMendeleev.
3. Modern formulation of the periodic law.
4. The value of the periodic law and the system of chemical elements of DIMendeleev.
5. Periodic system of chemical elements - a graphical reflection of the periodic law. The structure of the periodic system: periods, groups, subgroups.
6. Dependence of the properties of chemical elements on the structure of their atoms.
March 1 (according to the new style), 1869, is considered the date of the discovery of one of the most important laws of chemistry - the periodic law. In the middle of the XIX century. 63 chemical elements were known, and there was a need to classify them. Attempts at such a classification were made by many scientists (W. Odling and J. A. R. Newlands, J. B. A. Dumas and A. E. Chancourtua, I. V. Debereiner and L. Yu. Meyer), but only D. I. Mendeleev managed to see a certain pattern, arranging the elements in the order of increasing their atomic masses. This pattern has a periodic nature, so Mendeleev formulated the law he discovered as follows: the properties of the elements, as well as the forms and properties of their compounds, are in a periodic dependence on the value of the atomic mass of the element.
In the system of chemical elements proposed by Mendeleev, there were a number of contradictions that the author of the periodic law himself could not eliminate (argon-potassium, tellurium-iodine, cobalt-nickel). Only at the beginning of the 20th century, after the discovery of the structure of the atom, was the physical meaning of the periodic law explained and its modern formulation appeared: the properties of the elements, as well as the forms and properties of their compounds, are in a periodic dependence on the magnitude of the charge of the nuclei of their atoms.This formulation is confirmed by the presence of isotopes whose chemical properties are the same, although the atomic masses are different.
The Periodic Law is one of the fundamental laws of nature and the most important law of chemistry. With the discovery of this law, the modern stage in the development of chemical science begins. Although the physical meaning of the periodic law became clear only after the creation of the theory of the structure of the atom, this theory itself developed on the basis of the periodic law and the system of chemical elements. The law helps scientists to create new chemical elements and new compounds of elements, to obtain substances with the desired properties. Mendeleev himself predicted the existence of 12 elements that had not yet been discovered at that time, and determined their position in the periodic system. He described in detail the properties of three of these elements, and during the life of the scientist these elements were discovered (“ekabor” - gallium, “ekaaluminum” - scandium, “ekasilicon” - germanium). In addition, the periodic law is of great philosophical significance, confirming the most general laws of the development of nature.
Graphic reflection of the periodic law is the periodic system of chemical elements of Mendeleev. There are several forms of the periodic system (short, long, ladder (proposed by N. Bor), spiral). In Russia, the short form is the most widespread. The modern periodic system contains 110 chemical elements discovered to date, each of which occupies a certain place, has its own serial number and name. In the table, horizontal rows are distinguished - periods (1–3 are small, consist of one row; 4–6 are large, consist of two rows; the 7th period is incomplete). In addition to periods, vertical rows are distinguished - groups, each of which is divided into two subgroups (main - a and secondary - b). Secondary subgroups contain elements of only large periods, they all exhibit metallic properties. Elements of the same subgroup have the same structure of outer electron shells, which determines their similar chemical properties.
Period- this is a sequence of elements (from an alkali metal to an inert gas), the atoms of which have the same number of energy levels, equal to the number of the period.
Main subgroup is a vertical row of elements whose atoms have the same number of electrons in the outer energy level. This number is equal to the group number (except for hydrogen and helium).
All elements in the periodic system are divided into 4 electronic families ( s-, p-, d-,f-elements) depending on which sublevel in the element atom is filled last.
side subgroup is a vertical line d-elements that have the same total number of electrons per d-sublevel of the preexternal layer and s- sublevel of the outer layer. This number is usually equal to the group number.
The most important properties of chemical elements are metallicity and non-metallicity.
metallicity is the ability of the atoms of a chemical element to donate electrons. The quantitative characteristic of metallicity is the ionization energy.
Ionization energy of an atom- this is the amount of energy that is necessary to detach an electron from an atom of an element, i.e., to turn an atom into a cation. The lower the ionization energy, the easier the atom gives off an electron, the stronger the metallic properties of the element.
non-metallicity is the ability of atoms of a chemical element to attach electrons. The quantitative characteristic of non-metallicity is electron affinity.
electron affinity- this is the energy that is released when an electron is attached to a neutral atom, i.e., when an atom turns into an anion. The greater the affinity for an electron, the easier the atom attaches an electron, the stronger the non-metallic properties of the element.
A universal characteristic of metallicity and non-metallicity is the electronegativity (EO) of an element.
The EO of an element characterizes the ability of its atoms to attract electrons to themselves, which are involved in the formation of chemical bonds with other atoms in the molecule.
The more metallicity, the less EO.
The greater the non-metallicity, the greater the EO.
When determining the values of the relative EC on the Pauling scale, the EC of the lithium atom was taken as a unit (EC(Li) = 1); the most electronegative element is fluorine (EO(F) = 4).
In short periods from an alkali metal to an inert gas:
The charge of the nuclei of atoms increases;
The number of energy levels does not change;
The number of electrons in the outer level increases from 1 to 8;
The radius of the atoms decreases;
The strength of the bond between the electrons of the outer layer and the nucleus increases;
The ionization energy increases;
The electron affinity increases;
EO increases;
The metallicity of the elements decreases;
The non-metallicity of the elements increases.
Everything d-elements of this period are similar in their properties - they are all metals, have slightly different atomic radii and EC values, since they contain the same number of electrons at the outer level (for example, in the 4th period - except for Cr and Cu).
In the main subgroups from top to bottom:
The number of energy levels in an atom increases;
The number of electrons in the outer level is the same;
The radius of the atoms increases;
The strength of the bond between the electrons of the outer level and the nucleus decreases;
The ionization energy decreases;
The electron affinity decreases;
EO decreases;
The metallicity of the elements increases;
The non-metallicity of the elements decreases.
|
As a result of studying this topic, you will learn:
As a result of studying this topic, you will learn:
Study questions: |
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 open 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 the compounds of elements, are in a periodic dependence 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 due not to atomic mass, but to the electronic structure of atoms.
The nuclear charge is a characteristic that determines the electronic structure of atoms, and hence 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 in a periodic dependence on the serial number (on the magnitude of the charge of the nucleus of their atoms).
The expression of the Periodic Law is the periodic system of elements.
4.2. Periodic system of D. I. Mendeleev
The periodic system of elements of D. I. Mendeleev consists of seven periods, which are horizontal sequences of elements arranged in ascending 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 - large.
Each period (except the first) begins with alkali metal atoms (Li, Na, K, Rb, Cs, Fr) and ends with a noble gas (Ne, Ar, Kr, Xe, Rn) preceded by a typical non-metal. In periods from left to right, metallic properties gradually weaken and non-metallic properties increase, since with an increase in the positive charge of the nuclei of atoms, the number of electrons in the outer level increases.
In the first period, besides helium, there is only one element - hydrogen. It is conditionally placed in the IA or VIIA subgroup, 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, 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, Ca is followed by 10 transition elements (decade Sc - Zn), followed by the remaining 6 basic elements of the period (Ga - Kr). The fifth period is similarly constructed. concept transition element usually used to refer to any element with valence d- or f-electrons.
The sixth and seventh periods have double insertions of elements. The element Ba is followed by an intercalated decade of d-elements (La - Hg), and after the first transitional element La there are 14 f-elements - lanthanides(Se - Lu). After Hg 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). Recently, La and Ac have been classified as lanthanides and actinides, respectively. The lanthanides and actinides are placed separately at the bottom of the table.
Thus, each element in the periodic system occupies a strictly defined position, which is marked ordinal or atomic, number.
In the periodic system, eight groups (I - VIII) are located vertically, which in turn are divided into subgroups - main, or subgroups A and side, or subgroup B. Subgroup VIIIB is special, it contains triads elements that make up the families of iron (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 system. 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 this 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 system and electronic configurations of atoms
Since the nuclei of reacting atoms do not change during chemical reactions, the chemical properties of atoms depend on the structure of their electron shells.
The filling of electron layers and electron shells of atoms occurs in accordance with the Pauli principle and Hund's rule.
Pauli principle (Pauli prohibition)
Two electrons in an atom cannot have four identical quantum numbers (each atomic orbital can contain no more than two electrons).
The Pauli principle determines the maximum number of electrons that have a given principal quantum number n(i.e. located on a given electron layer): N n = 2n 2 . On the first electronic layer (energy level) there can be no more than 2 electrons, on the second - 8, on the third - 18, etc.
In the 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 equal to (2 l + 1). Each orbital is conventionally denoted by a cell, the electron located on it - 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 for an electron in a hydrogen atom in 1s orbitals
For both electrons of a helium atom n = 1, l = 0, m l= 0, m s = +1/2 and –1/2. Therefore, the electronic formula for helium is 1s 2 . The electron shell of helium is complete and very stable. Helium is a noble gas.
According to the Pauli principle, no two electrons with parallel spins can be in the same orbital. The third electron in the lithium atom occupies the 2s orbital. The electronic configuration of Li: 1s 2 2s 1, and for beryllium 1s 2 2s 2. Since the 2s orbital is filled, the fifth electron at the boron atom occupies the 2p orbital. At n= 2 side (orbital) quantum number l takes the values 0 and 1. When l = 0 (2s state) m l= 0, while l = 1 (2p is the state) m l can be equal to +1; 0; -one. The 2p state corresponds to three energy cells, fig. 4.2.
Rice. 4.2. The location of the electrons of the boron atom in orbitals
For a nitrogen atom (electronic configuration 1s 2 2s 2 2p 3 two electrons at the first level, five - at the second) the following two variants of the electronic structure are possible, fig. 4.3:
Rice. 4.3. Possible options for the arrangement of electrons of the nitrogen atom in orbitals
In the first scheme, Fig.4.3a, the total spin is 1/2 (+1/2 –1/2 +1/2), in the second (Fig.4.3b), the total spin is 3/2 (+1/2 + 1/2+1/2). The location of the spins is determined Hund's rule which reads: the energy levels are filled in such a way that the total spin is maximum.
In this way , of the two given schemes of 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 in the following way: first, one electron with identical spins, and then the second electron with opposite spins.
Starting with sodium, the third energy level with n = 3 is filled. 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 greatest bond with the nucleus. In 1961 V.M. Klechkovsky formulated a general position according to which the energy of electron orbitals increases in the order of increasing sum of the principal 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 ascending order of energy is roughly as follows:
1s< 2s < 2p < 3s < 3р < 4s ≈ 3d < 4p < 5s ≈ 4d < 5p < 6s ≈ 5d ≈ 4f < 6p.
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 Cr and Cu atoms by 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. Orbital distribution of electrons for a chromium atom
The physical reason for the "violation" of the order of filling is associated with the different penetrating power of the 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 layers of the charge kernels.
The electronic configurations of Mn, Fe, Co, Ni, Cu, and Zn atoms are represented 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 element 31 - gallium up to element 36 - krypton, the filling of the fourth layer (4p - shells) 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, electrons in excited states can be located in other atomic orbitals.
4.4. Types of chemical elements
All elements of the periodic system are divided into four types:
1. At atoms s-elements the s-shells of the outer layer (n) are filled. The s elements are hydrogen, helium, and the first two elements of each period.
2. At atoms p-elements electrons fill p-shells of the outer level (np). The p-elements include the last 6 elements of each period (except the first).
3. Do d-elements the d-shell of the second level outside (n-1) d is filled with electrons. These are elements of intercalated decades of large periods located between s- and p- elements.
4. Do f-elements filled with electrons f-sublevel of the third outside level (n-2) f . The family of f-elements includes lanthanides and actinides.
From the consideration of 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. Hence, s-elements are in all periods, p-elements are in the second and subsequent periods, d-elements are in the fourth and subsequent periods, and f-elements are 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 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 "dip" 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. For elements of secondary subgroups, the valence electrons are not only the outer, but also the 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. An exception is fluorine - its oxidation state is -1; Of the Group VIII elements, only Os, Ru, and Xe have an oxidation state of +8.
4.5. Periodicity of properties of atoms of elements
Such characteristics of atoms as their radius, ionization energy, electron affinity, electronegativity, 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 on the basis of 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 non-metals 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 conventional values.
From left to right, in periods, the value of the atomic radii of metals decreases, and the atomic radii of non-metals change in a complex way, since it depends on the nature of the chemical bond. In the second period, for example, the atomic radii first decrease and then increase, especially sharply when passing to a noble gas atom.
In the main subgroups, the atomic radii increase from top to bottom, as the number of electron layers increases.
The radius of the cation is less than the radius of the corresponding atom, and with an increase in the positive charge of the cation, its radius decreases. Conversely, 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 with decreasing negative and increasing positive ion radius. Such a decrease takes place, for example, in the series: O 2–, F –, Na +, Mg 2+, Al 3+.
Ionization energy is the energy required to detach an electron from an atom in the ground state. It is usually expressed in electronvolts (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 between the electron and the nucleus increases and the screening effect of the inner electron layers increases.
Table 4.1 shows the values of the ionization energies (energy of detachment of the first, second, etc. electrons) for some atoms.
In the second period, when passing from Li to Ne, the energy of detachment 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 following beryllium and nitrogen, respectively, its 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, therefore, in the next boron, an electron enters the p-orbital. This p-electron is less strongly 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 certain elements
Each p orbital of the nitrogen atom has one electron. At 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, so they have pronounced metallic properties, the highest ionization energy is in inert gases.
electron affinity is the energy released when an electron is attached to a neutral atom. Electron affinity, like ionization energy, is usually expressed in electron volts. Halogens have the highest electron affinity, while alkali metals have the lowest. Table 4.2 shows the values of electron affinity for atoms of some elements.
Table 4.2.
Electron affinity of atoms of some elements
Electronegativity- the ability of an atom in a molecule or ion to attract the valence electrons of other atoms. Electronegativity (EO) as a quantitative measure is an approximate value. About 20 electronegativity scales have been proposed, the most recognized of which was the scale developed by L. Pauling. On fig. 4.7 shows the values of EO according to Pauling.
Rice. 4.7. Electronegativity of the elements (according to Pauling)
Fluorine is the most electronegative of all elements on the Pauling scale. Its EO is taken equal to 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 acquires.
4.6. Acid-base properties of compounds; Kossel scheme
To explain the nature of the change in the acid-base properties of the compounds of the elements, Kossel (Germany) proposed using a simple scheme based on the assumption that a purely ionic bond exists in the molecules and that the Coulomb interaction takes place between the ions. The Kossel scheme describes the acid-base properties of compounds containing E-H and E-O-H bonds, depending on the charge of the nucleus and the radius of the element that forms them.
The Kossel scheme for two metal hydroxides, for example, LiOH and KOH, is shown in fig. 4.8.
Rice. 4.8. Kossel scheme for LiOH and KOH
As can be seen from the presented scheme, the radius of the Li + ion is less than the radius of the K + ion and the OH - group is more strongly bonded 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.
Similarly, one can analyze the Kossel scheme for the two bases CuOH and Cu(OH) 2 . Since the radius of the Cu 2+ ion is smaller and the charge is greater than that of the Cu + ion, the Cu 2+ ion will hold the OH - group more firmly. As a result, the Cu(OH) 2 base will be weaker than CuOH.
In this way, base strength increases as the cation radius increases and its positive charge decreases.
In the main subgroups, from top to bottom, the strength of the bases increases, since the radii of the element ions increase in this direction. In periods from left to right, there is a decrease in the radii of the ions of elements and an increase in their positive charge, therefore, in this direction, the strength of the bases decreases.
The Kossel scheme for two anoxic acids, for example, HCl and HI, is shown in fig. 4.9
Rice. 4.9. Kossel's scheme 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 the hydroiodic acid. In this way, the strength of anoxic acids increases with increasing negative ion radius.
The strength of oxygen-containing acids changes in the opposite way. It increases with decreasing ion radius and increasing its positive charge. On fig. 4.10 shows the Kossel scheme for two acids HClO and HClO 4 .
Rice. 4.10. Kossel scheme for HClO and HClO 4
The C1 7+ ion is strongly bound to the oxygen ion, so the proton will be more easily split off in the HClO 4 molecule. At the same time, the bond of the C1 + ion with the O 2- ion is less strong, and in the HC1O molecule the proton will be more strongly retained by the O 2- anion. As a result, HClO 4 will be a stronger acid than HClO.
The advantage of the Kossel scheme is that, using a simple model, it makes it possible to explain the nature of the change in the acid-base properties of compounds in a series of similar substances. However, this scheme is purely qualitative. It only allows one to compare the properties of compounds and does not make it possible to determine the acid-base properties of an arbitrarily chosen one compound. The disadvantage of this model is that it is based only on electrostatic concepts, while in nature there is no pure (100%) ionic bond.
4.7. Redox properties of elements and their compounds
A change in the redox properties of simple substances is easy to establish 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 oxidizing 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, while 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 system, 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. Oxidizing properties are the stronger, the greater the positive oxidation state of the element in the compound. Potassium permanganate and potassium dichromate show 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.
