Discovery of the periodic law and Mendeleev's periodic system. Introduction to General Chemistry

Data on the structure of the nucleus and the distribution of electrons in atoms make it possible to consider the periodic law and the periodic system of elements from fundamental physical positions. Based on modern concepts, the periodic law is formulated as follows:

Properties simple substances, as well as the forms and properties of compounds of elements are periodically dependent on the magnitude of the charge of the atomic nucleus (ordinal number).

Periodic table D.I. Mendeleev

Currently, more than 500 variants of the periodic table are known: this various shapes transmission of the periodic law.

The first version of the system of elements proposed by D.I. Mendeleev on March 1, 1869 was the so-called long form version. In this version, the periods were located on one line.

In the periodic system there are 7 periods horizontally, of which the first three are called small, and the rest - large. The first period contains 2 elements, the second and third - 8 each, the fourth and fifth - 18, the sixth - 32, the seventh (incomplete) - 21 elements. Each period, with the exception of the first, begins with an alkali metal and ends with a noble gas (the 7th period is unfinished).

All elements of the periodic table are numbered in the order in which they follow each other. The numbers of the elements are called atomic numbers or atomic numbers.

There are 10 rows in the system. Each small period consists of one row, each large period consists of two rows: even (upper) and odd (lower). In the even rows of large periods (fourth, sixth, eighth and tenth) there are only metals, and the properties of the elements in the row change slightly from left to right. In the odd rows of large periods (fifth, seventh and ninth), the properties of the elements in the row change from left to right, like typical elements.

The main feature by which elements of long periods are divided into two series is their oxidation state. Their identical values ​​are repeated twice in a period with increasing atomic masses elements. For example, in the fourth period, the oxidation states of elements from K to Mn change from +1 to +7, followed by the triad of Fe, Co, Ni (these are elements of the even series), after which the same increase in oxidation states is observed for elements from Cu to Br ( These are the elements of the odd row). We see the same in others long periods, excluding the seventh, which consists of one (even) row. The forms of combinations of elements are also repeated twice in large periods.

In the sixth period, following lanthanum, there are 14 elements with serial numbers 58-71, called lanthanides (the word “lanthanides” means like lanthanum, and “actinodes” means “like actinium”). They are sometimes called lanthanides and actinides, which means the following lanthanum, following actinium). The lanthanides are placed separately at the bottom of the table, and in the cell the asterisk indicates the sequence of their location in the system: La-Lu. The chemical properties of the lanthanides are very similar. For example, they are all reactive metals, react with water to form hydroxide and hydrogen. It follows from this that the lanthanides have a strong horizontal analogy.

In the seventh period, 14 elements with serial numbers 90-103 make up the actinide family. They are also placed separately - under the lanthanides, and in the corresponding cell two asterisks indicate the sequence of their location in the system: Ac-Lr. However, unlike the lanthanides, the horizontal analogy in the actinides is weakly expressed. They exhibit more different oxidation states in their compounds. For example, the oxidation state of actinium is +3, and uranium is +3, +4, +5 and +6. Studying the chemical properties of actinides is extremely difficult due to the instability of their nuclei.

There are eight groups arranged vertically in the periodic table (indicated by Roman numerals). The group number is associated with the degree of oxidation of the elements that they exhibit in compounds. Typically, the highest positive oxidation state of an element is equal to the group number. The exception is fluorine - its oxidation state is -1; copper, silver, gold exhibit oxidation states of +1, +2 and +3; Of the Group VIII elements, the oxidation state +8 is known only for osmium, ruthenium and xenon.

Group VIII contains noble gases. Previously it was believed that they were not capable of forming chemical compounds.

Each group is divided into two subgroups - the main and secondary ones, which in the periodic table is emphasized by the displacement of some to the right and others to the left. The main subgroup consists of typical elements (elements of the second and third periods) and those similar to them in chemical properties elements of long periods. The secondary subgroup consists only of metals - elements of long periods. Group VIII is different from the rest. In addition to the main helium subgroup, it contains three secondary subgroups: an iron subgroup, a cobalt subgroup and a nickel subgroup.

The chemical properties of the elements of the main and secondary subgroups differ significantly. For example, in group VII the main subgroup consists of non-metals F, CI, Br, I, At, and the secondary subgroup consists of metals Mn, Tc, Re. Thus, subgroups combine the elements that are most similar to each other.

All elements except helium, neon and argon form oxygen compounds; There are only 8 forms of oxygen compounds. In the periodic table they are often represented by general formulas, located under each group in order of increasing oxidation state of the elements: R 2 O, RO, R 2 O 3, RO 2, R 2 O 5, RO 3, R 2 O 7, RO 4, where R is an element of this group. The formulas of higher oxides apply to all elements of the group (major and minor), except in cases where the elements do not exhibit an oxidation state equal to the group number.

The elements of the main subgroups, starting from group IV, form gaseous hydrogen compounds, of which there are 4 forms. They are also represented by general formulas in the sequence RH 4, RH 3, RH 2, RH. The formulas of hydrogen compounds are located under the elements of the main subgroups and refer only to them.

The properties of elements in subgroups naturally change: from top to bottom, metallic properties increase and non-metallic properties weaken. Obviously, the metallic properties are most pronounced in francium, then in cesium; non-metallic - for fluorine, then - for oxygen.

The periodicity of the properties of elements can also be clearly traced by considering the electronic configurations of atoms.

The number of electrons located at the outer level in the atoms of elements, arranged in order of increasing atomic number, repeats periodically. The periodic change in the properties of elements with increasing atomic number is explained by a periodic change in the structure of their atoms, namely the number of electrons at their outer energy levels. Based on the number of energy levels in the electron shell of an atom, elements are divided into seven periods. The first period consists of atoms in which the electron shell consists of one energy level, in the second period - of two, in the third - of three, in the fourth - of four, etc. Each new period begins when a new energy level begins to be filled level.

In the periodic system, each period begins with elements whose atoms at the outer level have one electron - atoms of alkali metals - and ends with elements whose atoms at the outer level have 2 (in the first period) or 8 electrons (in all subsequent periods) - atoms of noble gases .

Next, we see that the outer electron shells are similar for atoms of elements (Li, Na, K, Rb, Cs); (Be, Mg, Ca, Sr); (F, Cl, Br, I); (He, Ne, Ar, Kr, Xe), etc. That is why each of the above groups of elements appears in a certain main subgroup of the periodic table: Li, Na, K, Rb, Cs in group I, F, Cl, Br, I - to VII, etc.

It is precisely because of the similarity in the structure of the electronic shells of atoms that their physical and chemical properties are similar.

Number main subgroups is determined by the maximum number of elements at the energy level and is equal to 8. The number of transition elements (elements side subgroups) is determined by the maximum number of electrons in the d-sublevel and is equal to 10 in each of the large periods.

Since in the periodic table of chemical elements D.I. Mendeleev, one of the side subgroups contains three transition elements that are similar in chemical properties (the so-called triads Fe-Co-Ni, Ru-Rh-Pd, Os-Ir-Pt), then the number of side subgroups, as well as the main ones, is equal to 8.

By analogy with transition elements, the number of lanthanides and actinides placed at the bottom of the periodic system in the form of independent rows is equal to the maximum number of electrons at the f-sublevel, i.e. 14.

The period begins with an element in whose atom there is one s-electron at the outer level: in the first period it is hydrogen, in the rest - alkali metals. The period ends with a noble gas: the first - with helium (1s 2), the remaining periods - with elements, the atoms of which at the external level have an electronic configuration ns 2 np 6 .

The first period contains two elements: hydrogen (Z = 1) and helium (Z = 2). The second period begins with the element lithium (Z = 3) and ends with neon (Z= 10). The second period has eight elements. The third period begins with sodium (Z = 11), the electronic configuration of which is 1s 2 2s 2 2p 6 3s 1. The filling of the third energy level began with it. It ends at the inert gas argon (Z = 18), the 3s and 3p sublevels of which are completely filled. Electronic formula of argon: 1s 2 2s 2 2p 6 3s 2 3p 6. Sodium is an analogue of lithium, argon is an analogue of neon. In the third period, as in the second, there are eight elements.

The fourth period begins with potassium (Z = 19), the electronic structure of which is expressed by the formula 1s 2 2s 2 2p 6 3s 2 3p64s 1. Its 19th electron occupied the 4s sublevel, the energy of which is lower than the energy of the 3d sublevel. The outer 4s electron gives the element properties similar to those of sodium. In calcium (Z = 20), the 4s sublevel is filled with two electrons: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2. From the scandium element (Z = 21), the filling of the 3d sublevel begins, since it is energetically more favorable than 4p -sublevel. Five orbitals of the 3d sublevel can be occupied by ten electrons, which is the case for atoms from scandium to zinc (Z = 30). Therefore, the electronic structure of Sc corresponds to the formula 1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 4s 2, and of zinc - 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2. In the atoms of subsequent elements up to the noble gas krypton (Z = 36) the 4p sublevel is being filled. The fourth period has 18 elements.

The fifth period contains elements from rubidium (Z = 37) to the noble gas xenon (Z = 54). The filling of their energy levels is the same as for elements of the fourth period: after Rb and Sr, ten elements from yttrium (Z= 39) up to cadmium (Z = 48) the 4d sublevel is filled, after which electrons occupy the 5p sublevel. In the fifth period, as in the fourth, there are 18 elements.

In the atoms of the elements of the sixth period of cesium (Z = 55) and barium (Z = 56) the 6s sublevel is filled. In lanthanum (Z = 57), one electron enters the 5d sublevel, after which the filling of this sublevel stops, and the 4f sublevel begins to be filled, the seven orbitals of which can be occupied by 14 electrons. This occurs in atoms of lanthanide elements with Z = 58 - 71. Since the deep 4f sublevel of the third level outside is filled in these elements, they have very similar chemical properties. From hafnium (Z = 72), the filling of the d sublevel resumes and ends at mercury (Z = 80), after which electrons fill the 6p sublevel. The filling of the level is completed at the noble gas radon (Z = 86). There are 32 elements in the sixth period.

The seventh period is unfinished. The filling of electronic levels with electrons is similar to the sixth period. After filling the 7s sublevel of France (Z = 87) and radium (Z = 88), an actinium electron enters the 6d sublevel, after which the 5f sublevel begins to be filled with 14 electrons. This occurs in atoms of actinide elements with Z = 90 - 103. After the 103rd element, the b d-sublevel is filled: in kurchatovium (Z = 104), nilsborium (Z = 105), elements Z = 106 and Z = 107. Actinides, like lanthanides, have many similar chemical properties.

Although the 3 d-sublevel is filled in after the 4s-sublevel, it is placed earlier in the formula, since all sublevels of a given level are written sequentially.

Depending on which sublevel is last filled with electrons, all elements are divided into four types (families).

1. s - Elements: the s-sublevel of the outer level is filled with electrons. These include the first two elements of each period.

2. p - Elements: the p-sublevel of the outer level is filled with electrons. These are the last 6 elements of each period (except the first and seventh).

3. d - Elements: the d sublevel of the second outside level is filled with electrons, and one or two electrons remain on the outer level (Pd has zero). These include elements of inserted decades of large periods located between the s- and p-elements (they are also called transition elements).

4. f - Elements: the f-sublevel of the third outside level is filled with electrons, and two electrons remain at the outer level. These are lanthanides and actinides.

In the periodic table there are 14 s-elements, 30 p-elements, 35 d-elements, 28 f-elements. Elements of the same type have a number of common chemical properties.

The periodic system of D.I. Mendeleev is a natural classification of chemical elements according to the electron structure of their atoms. The electronic structure of an atom, and therefore the properties of an element, is judged by the position of the element in the corresponding period and subgroup of the periodic system. The patterns of filling electronic levels explain the different number of elements in periods.

Thus, the strict periodicity of the arrangement of elements in D.I. Mendeleev’s periodic system of chemical elements is fully explained by the sequential nature of the filling of energy levels.

Conclusions:

The theory of atomic structure explains the periodic changes in the properties of elements. Increasing positive charges atomic nuclei from 1 to 107 determines the periodic repetition of the structure of the external energy level. And since the properties of elements mainly depend on the number of electrons in the outer level, they also repeat periodically. This is the physical meaning of the periodic law.

In short periods, with an increase in the positive charge of atomic nuclei, the number of electrons at the external level increases (from 1 to 2 - in the first period, and from 1 to 8 - in the second and third periods), which explains the change in the properties of elements: at the beginning of the period (except for the first period) there is an alkali metal, then the metallic properties gradually weaken and the non-metallic properties increase.

In large periods, as the charge of the nuclei increases, the filling of levels with electrons is more difficult, which also explains the more complex change in the properties of elements compared to elements of small periods. Thus, in even rows of large periods, with increasing charge, the number of electrons in the outer level remains constant and is equal to 2 or 1. Therefore, while the level next to the outer (second outer) is filled with electrons, the properties of the elements in these rows change extremely slowly. Only in odd rows, when the number of electrons in the outer level increases with increasing nuclear charge (from 1 to 8), the properties of the elements begin to change in the same way as those of typical ones.

In the light of the doctrine of the structure of atoms, the division of D.I. Mendeleev of all elements into seven periods. The period number corresponds to the number of energy levels of atoms filled with electrons. Therefore, s-elements are present 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 division of groups into subgroups, based on the difference in the filling of energy levels with electrons, is also easy to explain. For elements of the main subgroups, either s-sublevels (these are s-elements) or p-sublevels (these are p-elements) of external levels are filled. For elements of side subgroups, the (d-sublevel of the second outside level (these are d-elements) is filled. For lanthanides and actinides, the 4f- and 5f-sublevels are filled, respectively (these are f-elements). Thus, each subgroup combines elements whose atoms have similar structure of the outer electronic level. At the same time, the atoms of the elements of the main subgroups contain at the outer levels a number of electrons equal to the group number. The side subgroups include elements whose atoms have at the outer level two or one electron each.

Differences in structure also determine differences in the properties of elements of different subgroups of the same group. Thus, at the external level of the atoms of elements of the halogen subgroup there are seven electrons of the manganese subgroup - two electrons each. The former are typical metals, and the latter are metals.

But the elements of these subgroups also have general properties: entering chemical reactions, all of them (with the exception of fluorine F) can donate 7 electrons to form chemical bonds. In this case, the atoms of the manganese subgroup give up 2 electrons from the outer level and 5 electrons from the next level. Thus, for elements of side subgroups, the valence electrons are not only the outer, but also the penultimate (second outer) levels, which is the main difference in the properties of the elements of the main and side subgroups.

It also follows that the group number, as a rule, indicates the number of electrons that can participate in the formation of chemical bonds. This is the physical meaning of the group number.

So, the structure of atoms determines two patterns:

1) change in the properties of elements horizontally - in the period from left to right, metallic properties are weakened and non-metallic properties are enhanced;

2) change in the properties of elements vertically - in a subgroup, with increasing serial number, metallic properties increase and non-metallic properties weaken.

In this case, the element (and the cell of the system) is located at the intersection of the horizontal and vertical, which determines its properties. This helps to find and describe the properties of elements whose isotopes are obtained artificially.

This lesson examines the Periodic Law and the Periodic Table of Chemical Elements by D.I. Mendeleev in the light of the theory of atomic structure. The following concepts are explained: the modern formulation of the periodic law, the physical meaning of period and group numbers, the reasons for the periodicity of changes in the characteristics and properties of atoms of elements and their compounds using examples of small and large periods, main subgroups, the physical meaning of the periodic law, general characteristics an element and the properties of its compounds based on the element's position in the Periodic Table.

Topic: Structure of the atom. Periodic law

Lesson: Periodic Law and periodic table chemical elements D.I. Mendeleev

During the formation of the science of chemistry, scientists tried to systematize information about several dozen known by that time. This problem also fascinated D.I. Mendeleev. He looked for patterns and relationships that would cover all elements, and not just some of them. Mendeleev considered the most important characteristic of an element to be the mass of its atom. Having analyzed all the information known at that time about the chemical elements and arranging them in increasing order of their atomic masses, in 1869 he formulated the periodic law.

Statement of the law: the properties of chemical elements, simple substances, as well as the composition and properties of compounds are periodically dependent on the value of atomic masses.

At the time the periodic law was formulated, the structure of the atom and the existence of elementary particles were not yet known. It was also subsequently established that the properties of a substance do not depend on atomic masses, as Mendeleev assumed. Although, without this information, D.I. Mendeleev did not make a single mistake in his table.

After the discovery of Moseley, who established experimentally that the charge of the nucleus of an atom coincides with the serial number of the chemical element indicated by Mendeleev in his table, changes were made to the formulation of his law.

Modern wording of the law: the properties of chemical elements, simple substances, as well as the composition and properties of compounds are periodically dependent on the values ​​of the charges of atomic nuclei.

Rice. 1. The graphic expression of the periodic law is the Periodic Table of Chemical Elements by D. I. Mendeleev

Rice. 2. Let us consider the notation adopted in it using the example of rubidium

In each cell corresponding to an element, the following are presented: chemical symbol, name, serial number corresponding to the number of protons in the atom, relative atomic mass. The number of electrons in an atom corresponds to the number of protons. The number of neutrons in an atom can be found by the difference between the relative atomic mass and the number of protons, i.e. the atomic number.

N(n 0 ) = A r - Z

Quantity relative ordinal

neutrons atomic mass element number

For example, for the isotope of chlorine 35 Cl the number of neutrons is: 35-17= 18

The components of the periodic table are groups and periods.

The periodic table contains eight groups of elements. Each group consists of two subgroups: main and secondary. The main ones are indicated by the letter A, and the side ones - with a letter b. The main subgroup contains more elements than the secondary subgroup. The main subgroup contains s- and p-elements, the secondary subgroup contains d-elements.

Group- a column of the periodic table that combines chemical elements that are chemically similar due to similar electronic configurations of the valence layer. This fundamental principle construction of the periodic system. Let's consider this as an example of the elements of the first two groups.

Table 1

The table shows that the elements of the first group of the main subgroup have one valence electron. Elements of the second group of the main subgroup have two valence electrons.

Some main subgroups have their own special names:

Table 2

A string called a period is a sequence of elements arranged in order of increasing charge on their nuclei, starting with an alkali metal (or hydrogen) and ending with a noble gas.

Number period is equal number of electronic levels in an atom.

There are two main options for representing the periodic system: long-period, in which 18 groups are distinguished (Fig. 3) and short-period, in which there are 8 groups, but the concept of main and secondary subgroups is introduced (Fig. 1).

Homework

1. No. 3-5 (p. 22) Rudzitis G.E. Chemistry. Fundamentals of general chemistry. 11th grade: textbook for general education institutions: basic level / G.E. Rudzitis, F.G. Feldman. - 14th ed. - M.: Education, 2012.

2. Compare the electronic configuration of carbon and silicon atoms. What valence and oxidation states can they exhibit in chemical compounds? Give formulas for compounds of these elements with hydrogen. Give the formulas of their compounds with oxygen in the highest oxidation state.

3. Write the electronic formulas of the outer shells of the following elements: 14 Si, 15 P, 16 S, 17 Cl, 34 Se, 52 Te. Three elements from this series are chemical analogues (exhibit similar chemical properties). What are these elements?

Periodic law D.I. Mendeleev and the periodic table of chemical elements It has great importance in the development of chemistry. Let's plunge back to 1871, when chemistry professor D.I. Mendeleev, through numerous trials and errors, came to the conclusion that “... the properties of the elements, and therefore the properties of the simple and complex bodies, stand periodically depending on their atomic weight.” The periodicity of changes in the properties of elements arises due to the periodic repetition of the electronic configuration of the outer electron layer with an increase in the charge of the nucleus.

Modern formulation of the periodic law is this:

“the properties of chemical elements (i.e., the properties and form of the compounds they form) are periodically dependent on the charge of the nucleus of the atoms of the chemical elements.”

While teaching chemistry, Mendeleev understood that remembering the individual properties of each element caused difficulties for students. He began to look for ways to create system method to make it easier to remember element properties. The result was natural table, later it became known as periodic.

Our modern table is very similar to the periodic table. Let's take a closer look at it.

Mendeleev table

Mendeleev's periodic table consists of 8 groups and 7 periods.

The vertical columns of a table are called groups . Elements, within each group, have similar chemical and physical properties. This is explained by the fact that elements of the same group have similar electronic configurations of the outer layer, the number of electrons on which is equal to the group number. In this case, the group is divided into main and secondary subgroups.

IN Main subgroups includes elements whose valence electrons are located on the outer ns- and np-sublevels. IN Side subgroups includes elements whose valence electrons are located on the outer ns-sublevel and the inner (n - 1) d-sublevel (or (n - 2) f-sublevel).

All elements in periodic table , depending on which sublevel (s-, p-, d- or f-) valence electrons are classified into: s-elements (elements of the main subgroups of groups I and II), p-elements (elements of the main subgroups III - VII groups), d-elements (elements of side subgroups), f-elements (lanthanides, actinides).

The highest valency of an element (with the exception of O, F, elements of the copper subgroup and group eight) is equal to the number of the group in which it is found.

For elements of the main and secondary subgroups, the formulas of higher oxides (and their hydrates) are the same. In the main subgroups, the composition of hydrogen compounds is the same for the elements in this group. Solid hydrides form elements of the main subgroups of groups I - III, and groups IV - VII form gaseous hydrogen compounds. Hydrogen compounds of type EN 4 are more neutral compounds, EN 3 are bases, H 2 E and NE are acids.

The horizontal rows of a table are called periods. The elements in the periods differ from each other, but what they have in common is that the last electrons are at the same energy level ( principal quantum numbern- the same ).

The first period differs from the others in that there are only 2 elements: hydrogen H and helium He.

In the second period there are 8 elements (Li - Ne). Lithium Li, an alkali metal, begins the period, and the noble gas neon Ne closes it.

In the third period, just like in the second, there are 8 elements (Na - Ar). The period begins with the alkali metal sodium Na, and the noble gas argon Ar closes it.

The fourth period contains 18 elements (K - Kr) - Mendeleev designated it as the first large period. It also begins with the alkali metal Potassium and ends with the inert gas krypton Kr. The composition of large periods includes transition elements (Sc - Zn) - d- elements.

In the fifth period, similar to the fourth, there are 18 elements (Rb - Xe) and its structure is similar to the fourth. It also begins with the alkali metal rubidium Rb, and ends with the inert gas xenon Xe. The composition of large periods includes transition elements (Y - Cd) - d- elements.

The sixth period consists of 32 elements (Cs - Rn). Except 10 d-elements (La, Hf - Hg) it contains a row of 14 f-elements (lanthanides) - Ce - Lu

The seventh period is not over. It begins with Franc Fr, it can be assumed that it will contain, like the sixth period, 32 elements that have already been found (up to the element with Z = 118).

Interactive periodic table

If you look at periodic table and draw an imaginary line starting at boron and ending between polonium and astatine, then all metals will be to the left of the line, and non-metals to the right. Elements immediately adjacent to this line will have the properties of both metals and non-metals. They are called metalloids or semimetals. These are boron, silicon, germanium, arsenic, antimony, tellurium and polonium.

Periodic law

Mendeleev gave the following formulation of the Periodic Law: “properties 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.”
There are four main periodic patterns:

Octet rule states that all elements tend to gain or lose an electron in order to have the eight-electron configuration of the nearest noble gas. Because Since the outer s- and p-orbitals of noble gases are completely filled, they are the most stable elements.
Ionization energy is the amount of energy required to remove an electron from an atom. According to the octet rule, when moving across the periodic table from left to right, more energy is required to remove an electron. Therefore, elements on the left side of the table tend to lose an electron, and those on the right side tend to gain one. Inert gases have the highest ionization energy. The ionization energy decreases as you move down the group, because electrons at low energy levels have the ability to repel electrons at higher energy levels. This phenomenon is called shielding effect. Due to this effect, the outer electrons are less tightly bound to the nucleus. Moving along the period, the ionization energy smoothly increases from left to right.

Electron affinity– the change in energy when an atom of a substance in a gaseous state acquires an additional electron. As one moves down the group, the electron affinity becomes less negative due to the screening effect.

Electronegativity- a measure of how strongly it tends to attract electrons from another atom associated with it. Electronegativity increases when moving in periodic table from left to right and from bottom to top. It must be remembered that noble gases do not have electronegativity. Thus, the most electronegative element is fluorine.

Based on these concepts, let us consider how the properties of atoms and their compounds change in periodic table.

So, in a periodic dependence there are such properties of an atom that are associated with its electronic configuration: atomic radius, ionization energy, electronegativity.

Let us consider the change in the properties of atoms and their compounds depending on their position in periodic table of chemical elements.

The non-metallicity of the atom increases when moving in the periodic table left to right and bottom to top. Due to this the basic properties of the oxides decrease, A acid properties increase in the same order - when moving from left to right and from bottom to top. Moreover, the acidic properties of oxides are stronger, the higher the oxidation state of the element that forms it.

By period from left to right basic properties hydroxides weaken; in the main subgroups, from top to bottom, the strength of the foundations increases. Moreover, if a metal can form several hydroxides, then with an increase in the oxidation state of the metal, basic properties hydroxides weaken.

By period from left to right the strength of oxygen-containing acids increases. When moving from top to bottom within one group, the strength of oxygen-containing acids decreases. In this case, the strength of the acid increases with increasing oxidation state of the acid-forming element.

By period from left to right the strength of oxygen-free acids increases. When moving from top to bottom within one group, the strength of oxygen-free acids increases.

INTRODUCTION

Penza

Introduction

1. Periodic law of D. I. Mendeleev.

2. Structure of the periodic table.

3. Families of elements.

4. Sizes of atoms and ions.

5.Ionization energy is a quantitative measure of the reducing properties of atoms.

6. Electron affinity is a quantitative measure of the oxidizing properties of an atom.

7. Electronegativity of an atom is a quantitative measure of the redox properties of an element.

Conclusion.

Literature:

1. Korovin N.V. General chemistry. Textbook. – M.: graduate School, 1998. – p. 27 - 34.

Educational and material support:

1. Multimedia projector.

2. Short-period and long-period versions of tables of the periodic system D.I. Mendeleev.

3. Table of electronegativity of elements according to Pauling.

Purpose of the lesson:

Know: 1. Periodic law D.I. Mendeleev (formulation by D.I. Mendeleev and modern formulation). Structure of the periodic table. Serial number of the element, period, group, subgroup. S -, p-, d-, f - electronic properties of elements.

2.Atomic radii, ionization energy and electron affinity, electronegativity of elements, their changes by periods and groups.

Organizational and methodological instructions:

1.Check the availability of trainees and their readiness for classes, eliminate shortcomings.

2. Announce the topic and purpose of the lesson, educational issues, literature.

3. Justify the need to study this topic.

4. Consider educational questions using presentation frames and tables of the periodic table.

5. For each educational issue and at the end of the lesson sum up.

6.At the end of the lesson, give out a self-study task.

The fundamental law of nature and the theoretical basis of chemistry is the periodic law, discovered by D.I. Mendeleev in 1969 on the basis of deep knowledge in the field of chemistry and brilliant intuition. Later, the law received a theoretical interpretation based on models of atomic structure.

The first version of the periodic law was proposed by Mendeleev in 1869, and finally formulated in 1871.

Formulation of the periodic law 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.

In 1914, Moseley, studying the X-ray spectra of atoms, came to the conclusion that the atomic number of an element in the PS coincides with the charge of the nucleus of its atom.

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 nucleus of the element’s atoms.

Physical meaning of the periodic law(its connection with the structure of the atom):

The structure and properties of elements and their compounds are periodically dependent on the charge of the atomic nucleus and are determined by periodically repeating similar configurations of their atoms.

: 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.”

Story

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

Döbereiner triads and the first systems of elements

In the early 60s of the 19th century, several works appeared that immediately preceded the Periodic Law.

Spiral de Chancourtois

Newlands Octaves

Newlands Table (1866)

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

John Newlands was, of course, the first to give a series of elements arranged in order of increasing atomic masses, assign the corresponding atomic number to the chemical elements, and notice the systematic relationship between this order and the physicochemical properties of the elements. He wrote that in such a sequence the properties of elements are repeated, the equivalent weights (mass) of which differ by 7 units, or by a value that is a multiple of 7, i.e., as if the eighth element in order repeats the properties of the first, as in music the eighth note repeats first. Newlands tried to give this dependence, which actually occurs for light elements, a universal character. In his table, similar elements were located in horizontal rows, but in the same row there were often elements completely different in properties. In addition, Newlands was forced to place two elements in some cells; finally, the table did not contain any empty seats; As a result, the law of octaves was accepted with extreme skepticism.

Odling and Meyer tables

Manifestations of the periodic law in relation to electron affinity energy

The periodicity of the electron affinity energies of atoms is explained, naturally, by the same factors that were already noted when discussing ionization potentials (see definition of electron affinity energy).

They have the highest electron affinity p-elements of group VII. The lowest electron affinity is for atoms with the configuration s² ( , , ) and s²p 6 ( , ) or with half-filled p-orbitals ( , , ):

Manifestations of the periodic law regarding electronegativity

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

The periodicity of atomic electronegativity is important integral part periodic law and can easily be explained based on the immutable, although not entirely unambiguous, dependence of electronegativity values ​​on the corresponding values ​​of ionization energies and electron affinities.

In the periods there is a general tendency for electronegativity to increase, and in subgroups there is a decrease. The lowest electronegativity is for the s-elements of group I, the highest for the p-elements of group VII.

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

Rice. 4 Dependence of the orbital radii of atoms on the atomic number of the element.

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

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

Manifestations of the periodic law in relation to atomization energy

It should be emphasized that the oxidation state of an element, being a formal characteristic, does not provide an idea of ​​either the effective charges of the atoms of this element in the compound or the valency of the atoms, although the oxidation state is often called formal valency. Many elements are capable of exhibiting not one, but several different oxidation states. For example, for chlorine all oxidation states are known from −1 to +7, although the even ones are very unstable, and for manganese - from +2 to +7. The highest values ​​of the oxidation state change periodically depending on the atomic number of the element, but this periodicity is complex. In the simplest case, in the series of elements from an alkali metal to a noble gas, the highest oxidation state increases from +1 (F) to +8 (O4). In other cases, the highest oxidation state of the noble gas is less (+4 F 4) than for the preceding halogen (+7 O 4 −). Therefore, on the curve of the periodic dependence of the highest oxidation state on the atomic number of an element, the maxima fall either on the noble gas or on the halogen preceding it (the minima always on the alkali metal). The exception is the series - , in which neither the halogen () nor the noble gas () are known at all high degrees oxidation, and the middle member of the series, nitrogen, has the highest value of the highest oxidation state; therefore, in the series - the change in the highest oxidation state turns out to pass through a maximum. In general, the increase in the highest oxidation state in the series of elements from an alkali metal to a halogen or to a noble gas does not occur monotonically, mainly due to the manifestation of high oxidation states by transition metals. For example, the increase in the highest oxidation state in the series - from +1 to +8 is “complicated” by the fact that such high oxidation states as +6 (O 3), +7 (2 O 7), + are known for molybdenum, technetium and ruthenium 8(O4).

Manifestations of the periodic law in relation to oxidative potential

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

It is possible to detect some specific sequences in the changes in the oxidation potentials of simple substances. In particular, in the series of metals, during the transition from alkaline to the elements following them, a decrease in oxidation potentials occurs (+ (aq), etc. - hydrated cation):

This is easily explained by an increase in the ionization energy of atoms with an increase in the number of removed valence electrons. Therefore, on the curve of the dependence of the oxidation potentials of simple substances on the atomic number of the element, there are maxima corresponding to alkali metals. But it is not the only reason changes in the oxidation potentials of simple substances.

Internal and secondary periodicity

s- And R-elements

The general trends in the nature of changes in the values ​​of ionization energy of atoms, electron affinity energy of atoms, electronegativity, atomic and ionic radii, atomization energy of simple substances, oxidation state, oxidation potentials of simple substances depending on the atomic number of the element are discussed above. With a deeper study of these trends, one can find that the patterns in changes in the properties of elements in periods and groups are much more complex. In the nature of changes in the properties of elements over a period, internal periodicity is manifested, and in a group - secondary periodicity (discovered by E.V. Biron in 1915).

Thus, when passing from an s-element of group I to R-element of group VIII has internal maxima and minima on the atomic ionization energy curve and the curve of changes in their radii (see Fig. 1, 2, 4).

This indicates the internally periodic nature of the change in these properties over the period. An explanation of the noted patterns can be given using the concept of core shielding.

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

Rice. 5 Scheme of the structure of the last levels of beryllium, 9.32 eV (left) and boron, 8.29 eV (right)

This is explained by the fact that attraction to the nucleus 2p-electron of the boron atom is weakened due to the screening effect 2s-electrons.

It is clear that core shielding increases with increasing number of internal electronic layers. Therefore, in subgroups s- And R-elements there is a tendency to decrease the ionization energy of atoms (see Fig. 1).

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

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

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

Rice. 7 Secondary periodic dependence of the radii of atoms of outer p-orbitals on the atomic number

Rice. 8 Secondary periodic dependence of the first ionization energy of atoms on the atomic number

Rice. 9 Radial distribution of electron density in the sodium atom

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

The electron density concentration (the degree of electron penetration) at the same principal quantum number is greatest for s-electron, less - for R-electron, even less - for d-electron, etc. For example, with n = 3, the degree of penetration decreases in the sequence 3 s>3p>3d(see Fig. 10).

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

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

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

d-Elements

In the outer layer of atoms d-elements (with the exception of ) there are 1-2 electrons ( ns-state). The remaining valence electrons are located in (n-1) d-state, i.e. in the pre-external layer.

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

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

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

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

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

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

The periodic law is the basis of chemical systematics

see also

Notes

Literature

1. Akhmetov N. S. Current issues course in inorganic chemistry. - M.: Education, 1991. - 224 pp. - ISBN 5-09-002630-0
2. Korolkov D. V. Fundamentals of inorganic chemistry. - M.: Education, 1982. - 271 p.
3. Mendeleev D. I. Fundamentals of Chemistry, vol. 2. M.: Goskhimizdat, 1947. 389 p.
4. Mendeleev D.I.// Encyclopedic Dictionary of Brockhaus and Efron: In 86 volumes (82 volumes and 4 additional ones). - St. Petersburg. , 1890-1907.