home Botany book All characteristics change. Test “Patterns of change in the chemical properties of elements and their compounds by periods and groups. Patterns of changes in the properties of elements and their compounds by periods and groups

All characteristics change. Test “Patterns of change in the chemical properties of elements and their compounds by periods and groups. Patterns of changes in the properties of elements and their compounds by periods and groups

Explanatory note Thematic test "Patterns of change chemical properties elements and their compounds by periods and groups"designed to prepare students for the Unified State exam in chemistry. Target audience - 11th grade. Wording test tasks correspond to the demo version of the 2018 chemistry test and measurement materials.

The tasks are compiled by analogy with the tests published in the manual “USE. Chemistry: typical exam options: 30 options / ed. A.A. Kaverina", published by the publishing house " national education» (Moscow, 2017)

Patterns of changes in the chemical properties of elements and their compounds by periods and groups

1)Cl

2) K

3) Si

4) S

5) O

Of those in the series chemical elements select three elements that are in the Periodic Table of Chemical Elements D.I. Mendeleev are in the same period. Arrange the selected elements in decreasing order of their electronegativity.
Write in the answer field the numbers of the selected elements in the desired sequence.

Answer:

From the chemical elements indicated in the series, select three elements that are in the Periodic Table of Chemical Elements of D.I. Mendeleev are in the same group. Arrange the selected elements in ascending order of the acid properties of their hydrogen compounds.

From the chemical elements indicated in the series, select three elements that are in the Periodic Table of Chemical Elements of D.I. Mendeleev are in the same period. Arrange the selected elements in ascending order of the acid properties of their higher hydroxides.

From the chemical elements indicated in the series, select three elements that are in the Periodic Table of Chemical Elements of D.I. Mendeleev are in the same period. Arrange the selected elements in order of increasing number of outer electrons in the atoms of these elements.

From the chemical elements indicated in the series, select three elements that are in the Periodic Table of Chemical Elements of D.I. Mendeleev are in the same period. Arrange the selected elements in order of increasing oxidizing properties of their atoms.

From the chemical elements indicated in the series, select three elements that are in the Periodic Table of Chemical Elements of D.I. Mendeleev are in the same group. Arrange the selected elements in order of strengthening the main properties of the oxides they form.

From the chemical elements indicated in the row, select three metals. Arrange the selected elements in order of decreasing restorative properties.

From the chemical elements indicated in the series, select three elements that are in the Periodic Table of Chemical Elements of D.I. Mendeleev are in the same group.
Arrange these elements in order of increasing strength of attraction of valence electrons.

Answers

Question 1

Question 2

Question 3

Patterns of changes in the chemical properties of elements and their compounds by periods and groups

We list the patterns of changes in properties that are manifested within the periods:

— metallic properties decrease;

- non-metallic properties are enhanced;

— the degree of oxidation of elements in higher oxides increases from $+1$ to $+7$ ($+8$ for $Os$ and $Ru$);

— the degree of oxidation of elements in volatile hydrogen compounds increases from $-4$ to $-1$;

- oxides from basic through amphoteric are replaced by acid oxides;

- hydroxides from alkalis through amphoteric ones are replaced by acids.

D. I. Mendeleev in $ 1869 $ made a conclusion - he formulated the Periodic Law, which sounds like this:

The properties of chemical elements and the substances formed by them are in a periodic dependence on the relative atomic masses elements.

Systematizing chemical elements on the basis of their relative atomic masses, Mendeleev also paid great attention to the properties of the elements and the substances they form, distributing elements with similar properties into vertical columns - groups.

Sometimes, in violation of the regularity he revealed, Mendeleev put heavier elements with lower values of relative atomic masses. For example, he wrote in his table cobalt before nickel, tellurium before iodine, and when inert (noble) gases were discovered, argon before potassium. Mendeleev considered this arrangement necessary because otherwise these elements would fall into groups of elements dissimilar to them in properties, in particular, the alkali metal potassium would fall into the group of inert gases, and the inert gas argon into the group of alkali metals.

D. I. Mendeleev could not explain these exceptions from general rule, could not explain the reason for the periodicity of the properties of the elements and the substances formed by them. However, he foresaw that this reason lies in the complex structure of the atom, internal structure which had not been explored at the time.

In accordance with modern ideas about the structure of the atom, the basis for the classification of chemical elements is their charges atomic nuclei, and the modern formulation of the periodic law is:

The properties of chemical elements and the substances formed by them are in a periodic dependence on the charges of their atomic nuclei.

The periodicity in the change in the properties of the elements is explained by the periodic repetition in the structure of the external energy levels of their atoms. The number of energy levels total number the electrons located on them and the number of electrons at the outer level reflect the symbolism adopted in the Periodic system, i.e. disclose physical meaning period number, group number, and ordinal number of the element.

The structure of the atom also makes it possible to explain the reasons for the change in the metallic and non-metallic properties of elements in periods and groups.

The Periodic Law and the Periodic System of Chemical Elements of D. I. Mendeleev summarize information about chemical elements and the substances formed by them and explain the periodicity in the change in their properties and the reason for the similarity of the properties of elements of the same group. These two most important meanings of the Periodic Law and the Periodic System are complemented by another one, which is the ability to predict, i.e. predict, describe properties and indicate ways of discovering new chemical elements.

General characteristics of metals of the main subgroups of groups I ± III in connection with their position in the Periodic system of chemical elements of D. I. Mendeleev and the structural features of their atoms

Chemical elements - metals

Most of the chemical elements are classified as metals - $92$ of the $114$ known elements.

All metals except mercury are solids in their normal state and have a number of common properties.

Metals- These are malleable, ductile, ductile substances that have a metallic luster and are capable of conducting heat and electric current.

Atoms of metal elements donate electrons from the outer (and some of the outer) electron layer, turning into positive ions.

This property of metal atoms, as you know, is determined by the fact that they have relatively large radii and a small number of electrons (mainly from $1$ to $3$ on the outer layer).

The only exceptions are $6$ of metals: germanium, tin, and lead atoms have $4$ electrons on the outer layer, antimony and bismuth atoms have $5$, and polonium atoms have $6$.

Metal atoms are characterized by low electronegativity values (from $0.7$ to $1.9$) and exclusively reducing properties, i.e. the ability to donate electrons.

You already know that in the Periodic Table of Chemical Elements of D. I. Mendeleev, metals are below the boron-astatine diagonal, and also above it, in side subgroups. In periods and main subgroups, there are regularities known to you in changing the metallic, and hence the reducing properties of the atoms of the elements.

Chemical elements located near the boron-astatine diagonal ($Be, Al, Ti, Ge, Nb, Sb$) have dual properties: in some of their compounds they behave like metals, in others they exhibit the properties of non-metals.

In secondary subgroups, the reducing properties of metals most often decrease with increasing serial number.

This can be explained by the fact that the strength of the bond of valence electrons with the nucleus of the atoms of these metals is more affected by the value of the charge of the nucleus, and not by the radius of the atom. The value of the charge of the nucleus increases significantly, the attraction of electrons to the nucleus increases. In this case, although the radius of the atom increases, it is not as significant as that of the metals of the main subgroups.

Simple substances formed by chemical elements - metals, and complex metal-containing substances play an important role in the mineral and organic "life" of the Earth. Suffice it to recall that atoms (ions) of metal elements are integral part compounds that determine the metabolism in the human body, animals. For example, $76$ elements were found in human blood, of which only $14$ are not metals. In the human body, some elements - metals (calcium, potassium, sodium, magnesium) are present in in large numbers, i.e. are macronutrients. And such metals as chromium, manganese, iron, cobalt, copper, zinc, molybdenum are present in small quantities, i.e. this is trace elements.

Features of the structure of metals of the main subgroups of groups I-III.

alkali metals are metals of the main subgroup of group I. Their atoms at the outer energy level have one electron each. Alkali metals are strong reducing agents. Their reducing power and reactivity increase as the element's atomic number increases (i.e. from top to bottom in the Periodic Table). All of them have electronic conductivity. The strength of the bond between alkali metal atoms decreases with an increase in the atomic number of the element. Their melting and boiling points also decrease. Alkali metals interact with many simple substances - oxidizing agents. In reactions with water, they form water-soluble bases (alkalis).

Alkaline earth elements are called elements of the main subgroup of group II. The atoms of these elements contain two electrons at the outer energy level. They are reducing agents and have an oxidation state of $+2$. In this main subgroup, general patterns in the change in physical and chemical properties associated with an increase in the size of atoms in a group from top to bottom, the chemical bond between atoms also weakens. With an increase in the size of the ion, the acidic and the basic properties of oxides and hydroxides increase.

The main subgroup of group III consists of the elements boron, aluminum, gallium, indium and thallium. All elements refer to $p$-elements. At the outer energy level, they have three $(s^2p^1)$ electrons, which explains the similarity of properties. The oxidation state is $+3$. Within a group, as the nuclear charge increases, the metallic properties increase. Boron is a non-metal element, while aluminum already has metallic properties. All elements form oxides and hydroxides.

Characteristics of transition elements ± copper, zinc, chromium, iron according to their position in the Periodic system of chemical elements of D. I. Mendeleev and the structural features of their atoms

Most of the metal elements are in the side groups of the Periodic Table.

In the fourth period, the fourth electron layer appears at the potassium and calcium atoms, the $4s$-sublevel is filled, since it has a lower energy than the $3d$-sublevel. $K, Ca are s$-elements included in the main subgroups. For atoms from $Sc$ to $Zn$, the $3d$-sublevel is filled with electrons.

Consider what forces act on an electron that is added to an atom as the charge of the nucleus increases. On the one hand, the attraction of the atomic nucleus, which causes the electron to occupy the lowest free energy level. On the other hand, repulsion by already existing electrons. When there are $8$ electrons in the energy level ($s-$ and $p-$orbitals are occupied), their total repulsive effect is so strong that the next electron falls instead of the one located in energy below $d-$orbital into a higher $s-$ next level orbital. The electronic structure of the external energy levels of potassium is $...3d^(0)4s^1$, and that of calcium is $...3d^(0)4s^2$.

The subsequent addition of one more electron in scandium leads to the beginning of the filling of the $3d$-orbital instead of even higher-energy $4p$-orbitals. This turns out to be energetically more profitable. The filling of the $3d$ orbital ends with zinc, which has the electronic structure $1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)3d^(10)4s^2$. It should be noted that in the elements of copper and chromium, the phenomenon of "failure" of the electron is observed. The tenth $d$-electron of the copper atom moves to the third $3d$-sublevel.

The electronic formula of copper is $...3d^(10)4s^1$. A chromium atom in the fourth energy level ($s$-orbital) should have $2$ electrons. However, one of the two electrons goes to the third energy level, to the unfilled $d$-orbital, its electronic formula$...3d^(5)4s^1$.

Thus, in contrast to the elements of the main subgroups, where the atomic orbitals of the outer level are gradually filled with electrons, the $d$-orbitals of the penultimate energy level are filled in the elements of the secondary subgroups. Hence the name: $d$-elements.

All simple substances formed by elements of subgroups of the Periodic system are metals. Thanks to more atomic orbitals than the metal elements of the main subgroups, the atoms of $d$-elements form big number chemical bonds with each other and therefore create a stronger crystal lattice. It is stronger both mechanically and in relation to heating. Therefore, the metals of the secondary subgroups are the most durable and refractory among all metals.

It is known that if an atom has more than three valence electrons, then the element exhibits a variable valency. This provision applies to most $d$-elements. Their maximum valence, like the elements of the main subgroups, is equal to the group number (although there are exceptions). Elements with an equal number of valence electrons are included in the group under the same number $(Fe, Co, Ni)$.

For $d$-elements, the change in the properties of their oxides and hydroxides within one period when moving from left to right, i.e. with an increase in their valency, it proceeds from basic properties through amphoteric to acidic. For example, chromium has valencies $+2, +3, +6$; and its oxides: $CrO$ - basic, $Cr_(2)O_3$ - amphoteric, $CrO_3$ - acidic.

General characteristics of non-metals of the main subgroups of IV±VII groups in connection with their position in the Periodic system of chemical elements of D. I. Mendeleev and structural features of their atoms

Chemical elements - non-metals

The very first scientific classification of chemical elements was their division into metals and non-metals. This classification has not lost its significance at the present time.

non-metals These are chemical elements whose atoms are characterized by the ability to accept electrons before the completion of the outer layer due to the presence, as a rule, of four or more electrons on the outer electronic layer and the small radius of atoms compared to metal atoms.

This definition leaves aside the elements of group VIII of the main subgroup - inert, or noble, gases, the atoms of which have a complete outer electron layer. The electronic configuration of the atoms of these elements is such that they cannot be attributed to either metals or non-metals. They are those objects that separate elements into metals and non-metals, occupying a boundary position between them. Inert, or noble, gases (“nobility” is expressed in inertia) are sometimes referred to as non-metals, but formally, according to physical characteristics. These substances retain their gaseous state down to very low temperatures. Thus, helium does not go into a liquid state at $t°= -268.9 °C$.

The chemical inertness of these elements is relative. For xenon and krypton, compounds with fluorine and oxygen are known: $KrF_2, XeF_2, XeF_4$, etc. Undoubtedly, in the formation of these compounds, inert gases acted as reducing agents.

From the definition of non-metals, it follows that their atoms are characterized by high values of electronegativity. It varies from $2$ to $4$. Non-metals are elements of the main subgroups, mainly $p$-elements, with the exception of hydrogen - an s-element.

All non-metal elements (except hydrogen) occupy the upper right corner in the Periodic Table of Chemical Elements of D. I. Mendeleev, forming a triangle, the apex of which is fluorine $F$, and the base is the diagonal $B-At$.

However, special attention should be paid to the dual position of hydrogen in the Periodic system: in the main subgroups of groups I and VII. This is no coincidence. On the one hand, the hydrogen atom, like the atoms of alkali metals, has one electron on the outer (and only for it) electron layer (electronic configuration $1s^1$), which it is able to donate, exhibiting the properties of a reducing agent.

In most of its compounds, hydrogen, like alkali metals, exhibits an oxidation state of $+1$. But the release of an electron by a hydrogen atom is more difficult than that of alkali metal atoms. On the other hand, the hydrogen atom, like the halogen atoms, lacks one electron before the completion of the outer electron layer, so the hydrogen atom can accept one electron, exhibiting the properties of an oxidizing agent and the oxidation state characteristic of a halogen - $1$ in hydrides (compounds with metals, similar to compounds metals with halogens - halides). But the attachment of one electron to a hydrogen atom is more difficult than with halogens.

Properties of atoms of elements - non-metals

The atoms of non-metals are dominated by oxidizing properties, i.e. the ability to accept electrons. This ability characterizes the value of electronegativity, which naturally changes in periods and subgroups.

Fluorine is the strongest oxidizing agent, its atoms in chemical reactions are not able to donate electrons, i.e. exhibit restorative properties.

Configuration of the outer electron layer.

Other non-metals can exhibit reducing properties, although to a much weaker extent compared to metals; in periods and subgroups, their reducing ability changes in the reverse order compared to the oxidizing one.

Chemical elements-nonmetals only $16$! Quite a bit, considering that $114$ elements are known. Two non-metal elements make up $76%$ of the mass of the earth's crust. These are oxygen ($49%$) and silicon ($27%$). The atmosphere contains $0.03%$ of the mass of oxygen in earth's crust. Non-metals make up $98.5%$ of plant mass, $97.6%$ of human body mass. Nonmetals $C, H, O, N, S, P$ are organogens that form the most important organic matter living cell: proteins, fats, carbohydrates, nucleic acids. The composition of the air we breathe includes simple and complex substances, also formed by non-metal elements (oxygen $O_2$, nitrogen $N_2$, carbon dioxide$СО_2$, water vapor $Н_2О$, etc.).

Hydrogen is the main element of the universe. Many space objects (gas clouds, stars, including the Sun) are more than half made up of hydrogen. On Earth it, including the atmosphere, hydrosphere and lithosphere, is only $0.88%$. But this is by mass, and the atomic mass of hydrogen is very small. Therefore, its small content is only apparent, and out of every $100$ atoms on Earth, $17$ are hydrogen atoms.

The properties of elements and their compounds are determined: 1 - charges of atomic nuclei, 2 - atomic radii.

Small periods. Consider the change in some properties of elements and their compounds using the example of period II (see Table 3). In the second period, with an increase in the positive charge of the atomic nuclei, there is a sequential increase in the number of electrons at the outer level, which is the most distant from the atomic nucleus and therefore easily deformed, which leads to a rapid decrease in the atomic radius. This explains the rapid weakening of the metallic and reducing properties of elements, the strengthening of non-metallic and oxidizing properties, the increase in the acidic properties of oxides and hydroxides, and the decrease in basic properties. The period ends with a noble gas (Ne). In the third period, the properties of the elements and their compounds change in the same way as in the second, since the atoms of the elements of this period repeat the electronic structures of the atoms of the elements of the second period (3s- and 3p-sublevels)

Large periods (IV, V). In even rows of large periods (IV, V), starting from the third element, the number of electrons in the penultimate level increases sequentially, while the structure of the outer level remains unchanged. The penultimate level is located closer to the atomic nucleus and therefore deforms to a lesser extent. This leads to a slower decrease in the atomic radius. For example:

A consequence of a slow change in the radius of atoms and the same number of electrons at the outer level is also a slow decrease in the metallic and reducing properties of elements and their compounds. So, in an even row of IV period K - Mn - active metals Fe - Ni - metals of medium activity (compare with elements of period II, where the third element - boron - is already a non-metal).

And starting from group III of an odd series, the properties of elements and their compounds change in the same way as in small periods, since the external level begins to build up. Thus, the structure of the energy level is decisive in the properties of elements and their compounds. Each period under consideration also ends with a noble gas.

Having considered the change in some properties of elements and their compounds in periods, we can draw the following conclusions:

1. Each period begins with an alkali metal and ends with a noble gas.

2. The properties of elements and their compounds are periodically repeated because the structures of energy levels are periodically repeated. This is the physical meaning of the periodic law.

In the main subgroups, the number of energy levels increases, which leads to an increase in atomic radii. Therefore, in the main subgroups (from top to bottom), the electronegativity decreases, the megalithic and reducing properties of the elements increase, while the non-metallic and oxidizing properties decrease, the basic properties of oxides and hydroxides increase, and the acid properties decrease. For example, consider the main subgroup of group II.

Thus, the properties of an element and its compounds are intermediate between two elements adjacent to it in terms of period and subgroup.

According to the coordinates (period number and group number) of an element in the periodic system of D. I. Mendeleev, it is possible to determine the electronic structure of its atom, and, therefore, to foresee its main properties.

1. number of electronic levels in an atom defines period number The containing the corresponding element.

2. Total number of electrons, located in the s- and p-orbitals of the outer level (for elements of the main subgroups) and in the d-orbitals of the pre-outer and s-orbitals of the outer level (for elements of secondary subgroups; exceptions:

defines group number.

3. f-elements are located either in the side subgroup of group III (short-term variant), or between IIA- and IIIB-groups (long-term variant) - lanthanides(№ 57-70), actinides(№ 89-102).

4. atoms elements from different periods, but one subgroup have the same structure of the outer and pre-outer electronic levels and therefore have similar chemical properties.

5. element's maximum oxidation number coincides with the number of the group in which the element is located. The nature of the oxides and hydroxides formed by the element depends on oxidizing number of elements in them. Oxides and hydroxides in which the element is in the oxidation state:

The higher the degree of oxidation of the acid-forming element, the more pronounced acid properties oxides and hydroxides.

Therefore: oxides and hydroxides of elements of groups I-III are predominantly amphoteric. Oxides and hydroxides of elements of groups IV-VII are predominantly acidic (at the maximum degree of oxidation). Oxides and hydroxides of the same elements, but with a lower degree of oxidation, can be of a different nature.

6. Connections of elements with hydrogen can be subdivided into 3 major groups:

a) salt-like hydrides of active metals (LiH - , CaH - and etc.);

b) covalent hydrogen compounds of p-elements (B 2 H 6 , CH 4 , NH 3 , H 2 O, HF, etc.);

c) metal-like phases formed by d- and f-elements; the latter are usually non-stoichiometric compounds and it is often difficult to decide whether to refer to them as individual compounds or solid solutions.

Hydrogen compounds of elements of group IV (CH 4 -methane, SiH 4 - silane) do not interact with acids and bases, practically do not dissolve in water.

Hydrogen compounds of the elements of group V (NH 3 -ammonia) when dissolved in water form bases.

Hydrogen compounds of elements of groups VI and VII (H 2 S, HF) form acids when dissolved in water.

7. elements of the second period, in the atoms of which the 2nd electron layer is filled, are very different from all other elements. This is explained by the fact that the energy of electrons in the second layer is much lower than the energy of electrons in subsequent layers, and that the second layer cannot contain more than eight electrons.

8. d-elements of the same period differ less from each other than the elements of the main subgroups, in which the outer electronic layers are built up.

9. Differences in the properties of lanthanides, in the atoms of which the f-shell, which belongs to the third layer from the outside, is built up, are insignificant.

Every period(except for the first) begins with a typical metal and ends with a noble gas preceded by a typical non-metal.

Changing the properties of elements within a period:

1) weakening of metallic properties;

2) decrease in the radius of the atom;

3) strengthening of oxidizing properties;

4) the ionization energy increases;

5) electron affinity increases;

6) electronegativity increases;

7) acidic properties of oxides and hydroxides increase;

8) starting from group IV (for p-elements), the stability of hydrogen compounds increases and their acidic properties increase.

Changing the properties of elements within a group:

1) metallic properties increase;

2) the radius of the atom increases;

3) strengthening of reducing properties;

4) the ionization energy decreases;

5) electron affinity decreases;

6) electronegativity decreases;

7) the main properties of oxides and hydroxides increase;

8) starting from group IV (for p-elements), the stability of hydrogen compounds decreases, their acidic and oxidizing properties increase.

VALENCE- the ability of the atoms of elements to form chemical bonds. Quantitatively, valence is determined by the number of unpaired electrons.

In 1852, the English chemist Edward Frankland introduced the concept of connecting force. This property of atoms was later called valency.

valency is 2, because there are 2 unpaired electrons.

OXIDATION DEGREE- the conditional charge of the atom, which is calculated based on the assumption that the molecule consists only of ions.

Unlike valency, the oxidation state has a sign.

positive oxidation stateis equal to the number of drawn (given) electrons from a given atom. An atom can donate all unpaired electrons.

Negative oxidation stateis equal to the number of attracted (attached) electrons to a given atom; only non-metals show it. Atoms of non-metals attach such a number of electrons that is necessary to form a stable eight-electron configuration of the outer level.

For example: N -3 ; S-2; Cl-; C -4 .

The main pattern of this change is the strengthening of the metallic nature of the elements as Z increases. This pattern is especially pronounced in the IIIa-VIIa subgroups. For metals I A-III A-subgroups, an increase in chemical activity is observed. In the elements of IVA - VIIA subgroups, as Z increases, a weakening of the chemical activity of the elements is observed. For elements of b-subgroups, the change in chemical activity is more difficult.

Theory of the Periodic System was developed by N. Bohr and other scientists in the 20s. 20th century and is based on a real scheme for the formation of electronic configurations of atoms. According to this theory, as Z increases, the filling of electron shells and subshells in the atoms of elements included in the periods of the periodic system occurs in the following sequence:

Period numbers
1 2 3 4 5 6 7
1s 2s2p 3s3p 4s3d4p 5s4d5p 6s4f5d6p 7s5f6d7p

Based on the theory of the periodic system, one can give the following definition period: a period is a collection of elements beginning with the element with value n. equal to the period number, and l=0 (s-elements) and ending with an element with the same value n and l = 1 (p-elements) (see Atom). The exception is the first period containing only 1s elements. The number of elements in periods follows from the theory of the periodic system: 2, 8, 8. 18, 18, 32 ...

In the figure, the symbols of elements of each type (s-, p-, d- and f-elements) are depicted on a specific color background: s-elements - on red, p-elements - on orange, d-elements - on blue, f-elements - on green. Each cell contains the serial numbers and atomic masses of the elements, as well as the electronic configurations of the outer electron shells, which basically determine the chemical properties of the elements.

It follows from the theory of the periodic system that the a-subgroups include elements with and equal to the number of the period, and l = 0 and 1. The b-subgroups include those elements in whose atoms the shells that previously remained incomplete are completed. That is why the first, second and third periods do not contain elements of b-subgroups.

Structure of the Periodic Table of Chemical Elements closely related to the structure of atoms of chemical elements. As Z increases, similar types of configuration of the outer electron shells are periodically repeated. Namely, they determine the main features of the chemical behavior of elements. These features manifest themselves differently for the elements of the A-subgroups (s- and p-elements), for the elements of the b-subgroups (transitional d-elements), and for the elements of the f-families - lanthanides and actinides. A special case represent the elements of the first period - hydrogen and helium. Hydrogen is characterized by high chemical activity, because its only b-electron is easily split off. At the same time, the configuration of helium (1st) is very stable, which causes its complete chemical inactivity.

The elements of the A-subgroups are filled with outer electron shells (with n equal to the number of the period); therefore, the properties of these elements change noticeably as Z increases. Thus, in the second period, lithium (configuration 2s) - active metal, easily losing a single valence electron; beryllium (2s~) is also a metal, but less active due to the fact that its outer electrons are more firmly bound to the nucleus. Further, boron (2s "p) has a weakly pronounced metallic character, and all subsequent elements of the second period, in which the 2p subshell is built, are already non-metals. The eight-electron configuration of the outer electron shell of neon (2s ~ p ~) - an inert gas - is very durable.

Chemical properties of the elements of the second period are explained by the desire of their atoms to acquire the electronic configuration of the nearest inert gas (helium configuration - for elements from lithium to carbon or neon configuration - for elements from carbon to fluorine). This is why, for example, oxygen cannot exhibit a higher oxidation state equal to the group number: after all, it is easier for it to achieve the neon configuration by acquiring additional electrons. The same nature of the change in properties is manifested in the elements of the third period and in the s- and p-elements of all subsequent periods. At the same time, the weakening of the strength of the bond between the outer electrons and the nucleus in A-subgroups as Z increases manifests itself in the properties of the corresponding elements. Thus, for s-elements, there is a noticeable increase in chemical activity as Z increases, and for p-elements, an increase in metallic properties.

in atoms transitional d-elements previously unfinished shells with the value of the main quantum number and one less than the period number are completed. With some exceptions, the configuration of the outer electron shells of transition element atoms is ns. Therefore, all d-elements are metals, and that is why the changes in the properties of 1-elements as Z increases are not as sharp as we saw with s and p-elements. In higher oxidation states, d-elements show a certain similarity with p-elements of the corresponding groups of the periodic system.

The features of the properties of the elements of triads (VIII b-subgroup) are explained by the fact that the d-subshells are close to completion. This is why iron, cobalt, nickel and platinum metals, as a rule, are not inclined to give compounds of higher oxidation states. The only exceptions are ruthenium and osmium, which give the oxides RuO4 and OsO4. For elements of I- and II B-subgroups, the d-subshell actually turns out to be complete. Therefore, they exhibit oxidation states equal to the group number.

In the atoms of lanthanides and actinides (all of them are metals), the completion of previously incomplete electron shells with the value of the main quantum number and two units less than the period number takes place. In the atoms of these elements, the configuration of the outer electron shell (ns2) remains unchanged. At the same time, f-electrons do not actually affect the chemical properties. That's why the lanthanides are so similar.

For actinides, the situation is much more complicated. In the range of nuclear charges Z = 90 - 95, the electrons 6d and 5/ can take part in chemical interactions. And from this it follows that actinides exhibit a much wider range of oxidation states. For example, for neptunium, plutonium and americium, compounds are known where these elements act in a seven-valence state. Only for elements starting from curium (Z = 96) does the trivalent state become stable. Thus, the properties of the actinides differ significantly from those of the lanthanides, and therefore both families cannot be considered similar.

The actinide family ends with an element with Z = 103 (lawrencium). An evaluation of the chemical properties of kurchatovium (Z = 104) and nilsborium (Z = 105) shows that these elements should be analogues of hafnium and tantalum, respectively. Therefore, scientists believe that after the family of actinides in atoms, the systematic filling of the 6d subshell begins.

The finite number of elements that the periodic system covers is unknown. The problem of its upper limit is, perhaps, the main riddle of the periodic system. Most heavy element found in nature is plutonium (Z = 94). The reached limit of artificial nuclear fusion is an element with the serial number 107. It remains open question: will it be possible to get elements with large ordinal numbers, which ones and how many? It cannot yet be answered with any certainty.

Periodicity of properties of chemical elements

AT modern science the table of D. I. Mendeleev is called the periodic system of chemical elements, since the general patterns in the change in the properties of atoms, simple and complex substances, formed by chemical elements, are repeated in this system at certain intervals - periods. Thus, all chemical elements existing in the world are subject to a single, objectively acting in nature periodic law, the graphical representation of which is the periodic system of elements. This law and system bear the name of the great Russian chemist D. I. Mendeleev.

Periods- these are rows of elements arranged horizontally, with the same maximum value of the main quantum number of valence electrons. The period number corresponds to the number of energy levels in the element's atom. Periods consist of a certain number of elements: the first - from 2, the second and third - from 8, the fourth and fifth - from 18, the sixth period includes 32 elements. It depends on the number of electrons in the outer energy level. The seventh period is incomplete. All periods (with the exception of the first) begin with an alkali metal (s-element), and end with a noble gas. When a new energy level begins to fill, a new period begins. In a period with an increase in the ordinal number of a chemical element from left to right, metallic properties simple substances decrease, while non-metallic ones increase.

Metal properties is the ability of the atoms of an element to form chemical bond donate their electrons, and non-metallic properties are the ability of the atoms of an element to attach electrons to other atoms during the formation of a chemical bond. In metals, the outer s-sublevel is filled with electrons, which confirms the metallic properties of the atom. The non-metallic properties of simple substances manifest themselves during the formation and filling of the external p-sublevel with electrons. The non-metallic properties of the atom are enhanced in the process of filling the p-sublevel (from 1 to 5) with electrons. Atoms with a completely filled outer electron layer (ns 2 np 6) form a group noble gases which are chemically inert.

In short periods, with an increase in the positive charge of the nuclei of atoms, the number of electrons in the outer 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 the elements: at the beginning of the period (except for the first period) there is an alkali metal, then the metallic properties gradually weaken and non-metallic ones increase. For long periods As the nuclear charge increases, filling the levels with electrons becomes more difficult., which also explains a more complex change in the properties of elements compared to elements of small periods. So, in even rows of long periods, with increasing charge, the number of electrons in the outer level remains constant and is equal to 2 or 1. Therefore, while the electrons are filling the level following the outer (second from the outside) level, the properties of elements in even rows change extremely slowly. Only in odd rows, when the number of electrons in the outer level increases with the growth of the nuclear charge (from 1 to 8), do the properties of the elements begin to change in the same way as for typical ones.

Groups are vertical columns of elements with the same number of valence electrons equal to the group number. There is a division into main and secondary subgroups. The main subgroups consist of elements of small and large periods. The valence electrons of these elements are located at the outer ns- and np-sublevels. Side subgroups consist of elements of large periods. Their valence electrons are on the outer ns-sublevel and the inner (n - 1) d-sublevel (or (n - 2) f-sublevel). Depending on which sublevel (s-, p-, d- or f-) is filled with valence electrons, the elements are divided into:

1) s-elements - elements of the main subgroup of groups I and II;

2) p-elements - elements of the main subgroups of III-VII groups;

3) d-elements - elements of secondary subgroups;

4) f-elements - lanthanides, actinides.

Top down in the main subgroups, metallic properties are enhanced, while non-metallic properties are weakened. The elements of the main and secondary groups differ in properties. The group number indicates the highest valency of the element. The exceptions are oxygen, fluorine, elements of the copper subgroup and the eighth group. Common to the elements of the main and secondary subgroups are the formulas higher oxides(and their hydrates). In higher oxides and their hydrates of elements of groups I-III (with the exception of boron), basic properties predominate, from IV to VIII - acidic. For elements of the main subgroups, the formulas of hydrogen compounds are common. Elements of groups I-III form solid substances - hydrides, since the oxidation state of hydrogen is -1. Elements IV-VII groups - gaseous. Hydrogen compounds of the elements of the main subgroups of group IV (EN 4) are neutral, group V (EN3) are bases, groups VI and VII (H 2 E and NE) are acids.

Radii of atoms, their periodic changes in the system of chemical elements

The radius of an atom with an increase in the charges of the nuclei of atoms in a period decreases, since the attraction of the electron shells by the nucleus is enhanced. There is a kind of "compression". From lithium to neon, the charge of the nucleus gradually increases (from 3 to 10), which causes an increase in the forces of attraction of electrons to the nucleus, the size of atoms decreases. Therefore, at the beginning of the period, there are elements with a small number of electrons in the outer electron layer and a large atomic radius. Electrons that are farther from the nucleus are easily detached from it, which is typical for metal elements.

In the same group, with increasing period number, atomic radii increase, since an increase in the charge of an atom has the opposite effect. From the point of view of the theory of the structure of atoms, the belonging of elements to metals or non-metals is determined by the ability of their atoms to give or add electrons. Metal atoms donate electrons relatively easily and cannot add them to complete the construction of their outer electronic layer.

D. I. Mendeleev in 1869 formulated the periodic law, which sounds like this: the properties of chemical elements and the substances formed by them are in a periodic dependence on the relative atomic masses of the elements. Systematizing chemical elements on the basis of their relative atomic masses, Mendeleev also paid great attention to the properties of the elements and the substances they formed, distributing elements with similar properties into vertical columns - groups. In accordance with modern ideas about the structure of the atom, the basis for the classification of chemical elements is the charges of their atomic nuclei, and the modern formulation of the periodic law is as follows: the properties of chemical elements and the substances they form are in a periodic dependence on the charges of their atomic nuclei. The periodicity in the change in the properties of the elements is explained by the periodic repetition in the structure of the external energy levels of their atoms. It is the number of energy levels, the total number of electrons located on them, and the number of electrons at the outer level that reflect the symbolism adopted in the periodic system.

a) Patterns associated with metallic and non-metallic properties of elements.

b) Patterns associated with redox properties. Changes in the electronegativity of elements.

The above reasons explain why FROM LEFT TO RIGHT OXIDATIVE properties, and when moving TOP DOWN - RECOVERY element properties.
The latter regularity extends even to such unusual elements as inert gases. In the "heavy" noble gases of krypton and xenon, which are in the lower part of the group, it is possible to "select" electrons and obtain their compounds with strong oxidizing agents (fluorine and oxygen), but for the "light" helium, neon and argon this cannot be done.
In the upper right corner of the table is the most active non-metal oxidizer, fluorine (F), and in the lower left corner, the most active reducing metal, cesium (Cs). The element francium (Fr) should be an even more active reducing agent, but its chemical properties are extremely difficult to study due to its rapid radioactive decay.
For the same reason as the oxidizing properties of the elements, their ELECTRICITY INCREASES too FROM LEFT TO RIGHT, reaching a maximum for halogens. Not last role this is played by the degree of completion of the valence shell, its proximity to the octet.
When moving TOP DOWN by groups ELECTRICITY DECREASES. This is due to an increase in the number of electron shells, on the last of which electrons are attracted to the nucleus more and more weakly.
c) Regularities related to the size of atoms.
Atom sizes (ATOMIC RADIUS) when moving FROM LEFT TO RIGHT along the period DECREASE. Electrons are attracted more and more to the nucleus as the charge of the nucleus increases. Even an increase in the number of electrons in the outer shell (for example, in fluorine compared to oxygen) does not lead to an increase in the size of the atom. Conversely, the size of a fluorine atom is smaller than that of an oxygen atom.
When moving FROM TOP DOWN ATOMIC RADIUS elements GROW, because more electron shells are filled.

d) Patterns associated with the valency of elements.

elements of the same SUB-GROUPS have a similar configuration of outer electron shells and therefore the same valence in compounds with other elements.
s-elements have valences that match their group number.
p-Elements have the highest possible valence for them, equal to the group number. In addition, they can have a valency equal to the difference between the number 8 (octet) and their group number (the number of electrons in the outer shell).
The d-elements exhibit many different valences that cannot be accurately predicted from the group number.
Not only the elements, but also many of their compounds—oxides, hydrides, compounds with halogens—display periodicity. For each GROUPS elements, you can write the formulas of the compounds, which are periodically "repeated" (that is, they can be written as a generalized formula).

So, let's summarize the patterns of changes in properties, manifested within the periods:

Change of some characteristics of elements in periods from left to right:

the radius of the atoms decreases;
the electronegativity of the elements increases;
the number of valence electrons increases from 1 to 8 (equal to the group number);
the highest oxidation state increases (equal to the group number);
the number of electron layers of atoms does not change;
metallic properties are reduced;
the non-metallic properties of the elements are increased.

Changing some characteristics of elements in a group from top to bottom:

the charge of the nuclei of atoms increases;
the radius of the atoms increases;
the number of energy levels (electronic layers) of atoms increases (equal to the period number);
the number of electrons on the outer layer of atoms is the same (equal to the group number);
the strength of the bond between the electrons of the outer layer and the nucleus decreases;
electronegativity decreases;
the metallicity of the elements increases;
the non-metallicity of the elements decreases.

Z is a serial number, is equal to the number protons; R is the radius of the atom; EO - electronegativity; Shaft e - the number of valence electrons; OK. St. — oxidizing properties; Sun. St. - restorative properties; En. ur. — energy levels; Me - metallic properties; NeMe - non-metallic properties; BCO - the highest degree of oxidation

Reference material for passing the test:

periodic table

Solubility table