Modern formulation periodic law : 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 nuclei of their atoms ( serial number).

    Periodic properties are, for example, atomic radius, ionization energy, electron affinity, electronegativity of the atom, as well as some physical properties elements and compounds (melting and boiling points, electrical conductivity, etc.).

    The expression of the Periodic Law is

     periodic table of elements .

    The most common short form option periodic table, in which the elements are divided into 7 periods and 8 groups.

    Currently, the nuclei of atoms of elements up to number 118 have been obtained. The name of the element with serial number 104 is rutherfordium (Rf), 105 – dubnium (Db), 106 – seaborgium (Sg), 107 – bohrium (Bh), 108 – hassium (Hs ), 109 – meitnerium ( Mt), 110 - darmstadtium (Ds), 111 - roentgenium (Rg), 112 - copernicium (Cn).
    On October 24, 2012, in Moscow, at the Central House of Scientists of the Russian Academy of Sciences, a solemn ceremony was held to assign the name “flerovium” (Fl) to the 114th element, and “livermorium” (Lv) to the 116th element.

    Periods 1, 2, 3, 4, 5, 6 contain 2, 8, 8, 18, 18, 32 elements, respectively. The seventh period is not completed. Periods 1, 2 and 3 are called small, the rest - big.

    In periods from left to right, metallic properties gradually weaken and non-metallic properties increase, since with an increase in the positive charge of atomic nuclei, the number of electrons in the outer electronic layer increases and a decrease in atomic radii is observed.

    At the bottom of the table are 14 lanthanides and 14 actinides. Recently, lanthanum and actinium have been classified as lanthanides and actinides, respectively.

    Groups are divided into subgroups - main ones, or subgroups A and side effects, or subgroup B. Subgroup VIII B – special, it contains triads elements that make up the families of iron (Fe, Co, Ni) and platinum metals (Ru, Rh, Pd, Os, Ir, Pt).

    From top to bottom in the main subgroups, metallic properties increase and non-metallic properties weaken.

    The group number usually indicates the number of electrons that can participate in the formation of chemical bonds. This is physical meaning group numbers. Elements of side subgroups have valence electrons not only in the outer layers, but also in the penultimate layers. This is the main difference in the properties of the elements of the main and secondary subgroups.

    Periodic table and electronic formulas of atoms

    To predict and explain the properties of elements, you must be able to write the electronic formula of an atom.

    In an atom located in ground condition, each electron occupies a vacant orbital with the lowest energy. The energy state is determined primarily by temperature. The temperature on the surface of our planet is such that the atoms are in the ground state. At high temperatures, other states of atoms, which are called excited.

    Arrangement sequence energy levels in order of increasing energy is known from the results of solving the Schrödinger equation:

    1s< 2s < 2p < 3s < Зр < 4s 3d < 4p < 5s 4d < 5p < 6s 5d 4f < 6p.

    Let's consider the electronic configurations of atoms of some elements fourth period(Fig. 6.1).

    Rice. 6.1. Distribution of electrons over the orbitals of some elements of the fourth period

    It should be noted that there are some features in electronic structure atoms of elements of the fourth period: for atoms Cr and C u by 4 s-shell contains not two electrons, but one, i.e. "failure" external s -electron to the previous one d-shell.

    Electronic formulas of 24 Cr and 29 Cu atoms can be represented as follows:

    24 Cr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 1,

    29 Cu 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 1 .

    Physical reason“Violation” of the filling order is associated with different penetration abilities of electrons into the inner layers, as well as the special stability of the electronic configurations d 5 and d 10, f 7 and f 14.

    All elements are divided into four types

         :

    1. In atoms s-elements filled in s - outer layer shell ns . These are the first two elements of each period.

    2. At atoms p-elements electrons fill the p-shells of the outer np level . These include the last 6 elements of each period (except the first and seventh).

    3. U d-elements filled with electrons d -sublevel of the second outside level ( n-1)d . These are elements of intercalary decades of large periods located between s- and p-elements.

    4. U f-elements filled with electrons f -sublevel of the third outside level ( n-2)f . These are lanthanides and actinides.

    Changes in the acid-base properties of element compounds by groups and periods of the periodic system
    (Kossel diagram)

    To explain the nature of the change in the acid-base properties of compounds of elements, Kossel (Germany, 1923) proposed using a simple scheme based on the assumption that there is pure ionic bond and there is a Coulomb interaction between the ions. The Kossel scheme describes the acid-base properties of compounds containing E–H and E–O–H bonds, depending on the charge of the nucleus and the radius of the element forming them.

    Kossel diagram for two metal hydroxides (for LiOH and KOH molecules ) is shown in Fig. 6.2. As can be seen from the presented diagram, the radius of the Li ion + less than the ion radius K+ and OH The - - group is bonded more tightly to the lithium ion than to the potassium ion. As a result, KOH will be easier to dissociate in solution and the basic properties of potassium hydroxide will be more pronounced.

    Rice. 6.2. Kossel diagram for LiOH and KOH molecules

    In a similar way, you can analyze the Kossel scheme for two bases CuOH and Cu(OH) 2 . Since the radius of the Cu ion 2+ less, and the charge is greater than that of the ion Cu+, OH - - the group will be held more firmly by the Cu 2+ ion .
    As a result, the base     Cu(OH)2 will be weaker than CuOH.

    Thus, the strength of bases increases as the radius of the cation increases and its positive charge decreases .

    Kossel diagram for two oxygen-free acids HCl and HI shown in Fig. 6.3.

    Rice. 6.3. Kossel diagram for HCl and HI molecules

    Since the radius of the chloride ion is smaller than that of the iodide ion, the H+ ion more strongly bound to the anion in the hydrochloric acid molecule, which will be weaker than hydroiodic acid. Thus, the strength of anoxic acids increases with increasing radius of the negative ion.

    Force oxygen-containing acids changes in the opposite way. It increases as the radius of the ion decreases and its positive charge increases. In Fig. Figure 6.4 shows the Kossel diagram for two acids HClO and HClO 4.

    Rice. 6.4. Kossel diagram for HClO and HClO 4

    Ion C1 7+ is firmly bound to the oxygen ion, so the proton will be more easily split off in the HC1O molecule 4 . At the same time, the bond of the C1 ion+ with O 2- ion less strong, and in the HC1O molecule the proton will be more strongly retained by the O anion 2- . As a result, HClO 4 is a stronger acid than HClO.

    Thus, An increase in the oxidation state of an element and a decrease in the radius of the element’s ion increase the acidic nature of the substance. On the contrary, a decrease in the oxidation state and an increase in the ion radius enhance the basic properties of substances.

    Examples of problem solving

         Compose electronic formulas of the zirconium atom and ions     O 2– , Al 3+ , Zn 2+ . Determine what type of elements the Zr, O, Zn, Al atoms belong to.

      40 Zr 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 2 5s 2,

      O 2– 1s 2 2s 2 2p 6,

      Zn 2+ 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 ,

      Al 3+ 1s 2 2s 2 2p 6 ,

    Zr – d-element, O – p-element, Zn – d-element, Al – p-element.

    Arrange the atoms of the elements in order of increasing their ionization energy: K, Mg, Be, Ca. Justify the answer.

    Solution. Ionization energy– the energy required to remove an electron from an atom in the ground state. In the period from left to right, the ionization energy increases with increasing nuclear charge; in the main subgroups from top to bottom it decreases as the distance from the electron to the nucleus increases.

    Thus, the ionization energy of atoms of these elements increases in the series K, Ca, Mg, Be.

     Arrange atoms and ions in increasing order of their radii: Ca 2+, Ar, Cl –, K +, S 2– . Justify the answer.

    Solution. For ions containing the same number of electrons (isoelectronic ions), the radius of the ion will increase as its positive charge decreases and its negative charge increases. Consequently, the radius increases in the order Ca 2+, K +, Ar, Cl –, S 2–.

    Determine how the radii of ions and atoms change in the series Li + , Na + , K + , Rb + , Cs + and Na, Mg, Al, Si, P, S.

    Solution. In the series Li + , Na + , K + , Rb + , Cs + the radius of ions increases as the number of electronic layers of ions of the same sign with a similar electronic structure increases.

    In the series Na, Mg, Al, Si, P, S, the radius of atoms decreases, since with the same number of electron layers in the atoms, the charge of the nucleus increases, and, therefore, the attraction of electrons by the nucleus increases.

     Compare the strength of acids H 2 SO 3 and H 2 SeO 3 and bases Fe(OH) 2 and Fe(OH) 3.

    Solution. According to the Kossel scheme H 2 SO 3 more strong acid, than H 2 SeO 3 , since the ion radius SE 4+ greater than the ion radius S 4+, which means the S 4+ – O 2– bond is stronger than bond Se 4+ – O 2– .

    According to the Kossel scheme Fe(OH)

     2 stronger base since the radius of the Fe ion 2+ more than Fe ion 3+ . In addition, the charge of the Fe ion 3+ greater than that of Fe ion 2+ . As a result, the Fe bond 3+ – О 2– is stronger than Fe 2+ – O 2– and ION – easier to split off in a molecule Fe(OH)2.

    Problems to solve independently

    6.1.Create electronic formulas for elements with a nuclear charge of +19, +47, +33 and those in the ground state. Indicate what type of elements they belong to. What oxidation states are characteristic of an element with a nuclear charge of +33?

    6.2.Write the electronic formula of the Cl ion – .

The main pattern of this change is the strengthening of the metallic character of the elements as Z increases. This pattern is especially clearly manifested in the IIIa-VIIa subgroups. For metals I A-III A-subgroups there is an increase chemical activity. For elements of IVA - VIIA subgroups, as Z increases, a weakening of the chemical activity of the elements is observed. For b-subgroup elements, the change in chemical activity is more complex.

Theory of the periodic table was developed by N. Bohr and other scientists in the 20s. XX 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 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, the following definition of a period can be given: a period is a set of elements, starting with an 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, which contains only 1s elements. From the theory of the periodic system, the number of elements in the periods follows: 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 shows the atomic numbers and atomic masses of the elements, as well as the electronic configurations of the outer electron shells, which mainly determine Chemical properties elements.

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

Structure of the periodic table chemical elements is closely related to the structure of atoms of chemical elements. As Z increases, similar types of configuration of the outer electron shells periodically repeat. Namely, they determine the main features of the chemical behavior of elements. These features manifest themselves differently for elements of the A-subgroups (s- and p-elements), for elements of the b-subgroups (transition d-elements) and elements of the f-families - lanthanides and actinides. A special case is represented by the elements of the first period - hydrogen and helium. Hydrogen is highly reactive because its single b electron is easily removed. At the same time, the configuration of helium (1st) is very stable, which determines its complete chemical inactivity.


For elements of A-subgroups, the outer electron shells are filled (with n equal to the period number); therefore, the properties of these elements change noticeably as Z increases. Thus, in the second period, lithium (2s configuration) is an active metal that easily loses its only valence electron; beryllium (2s~) is also a metal, but less active due to the fact that its outer electrons are more tightly bound to the nucleus. Further, boron (23"p) has a weakly expressed 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 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 its group number: it is easier for it to achieve the neon configuration by acquiring additional electrons. The same nature of changes 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 outer electrons and the nucleus in A-subgroups as Z increases is manifested 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 there is an increase in metallic properties.

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

The peculiarities 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, do not tend to produce compounds in higher oxidation states. The only exceptions are ruthenium and osmium, which give the oxides RuO4 and OsO4. For elements of the I- and II B-subgroups, the d-subshell is actually complete. Therefore, they exhibit oxidation states equal to the group number.

In the atoms of lanthanides and actinides (all of them are metals), previously incomplete electron shells are completed with a value of the principal quantum number and two units less than the period number. In the atoms of these elements, the configuration of the outer electron shell (ns2) remains unchanged. At the same time, f electrons actually have no effect on chemical properties. This is why the lanthanides are so similar.

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

The actinide family ends with the element with Z = 103 (lawrencium). An assessment 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 actinide family in atoms, the systematic filling of the 6d subshell begins.

Final number elements that the periodic table covers is unknown. The problem of its upper limit is perhaps the main mystery of the periodic table. The heaviest element that has been discovered in nature is plutonium (Z = 94). The limit of artificial nuclear fusion has been reached - element with atomic number 107. Remains open question: will it be possible to obtain elements with large serial numbers, which ones and how many? This cannot yet be answered with any certainty.

Periodicity of properties of chemical elements

IN modern science D. I. Mendeleev’s table is called the periodic system of chemical elements, since there are general patterns in changes 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 periodic law objectively operating in nature, the graphic representation of which is the periodic system of elements. This law and system are named after the great Russian chemist D.I. Mendeleev.

Periods- these are rows of elements located 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 an element's atom. The periods consist of a certain number of elements: the first - of 2, the second and third - of 8, the fourth and fifth - of 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, it begins new period. In a period with an increase in the serial number of a chemical element from left to right, the metallic properties of simple substances decrease, and non-metallic ones increase.

Metallic properties is the ability of the atoms of an element to form chemical bond give up their electrons, and nonmetallic properties are the ability of the atoms of an element to attach electrons of other atoms when forming a chemical bond. In metals, the outer s-sublevel is filled with electrons, which confirms the metallic properties of the atom. The nonmetallic properties of simple substances manifest themselves during the formation and filling of the outer p-sublevel with electrons. The nonmetallic properties of the atom are enhanced by 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, as the positive charge of atomic nuclei increases, 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 alkali metal, then the metallic properties gradually weaken and non-metallic properties increase. In long periods As the charge of nuclei increases, filling levels with electrons becomes 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 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 level next to the outer (second from the outside) is filled with electrons, the properties of the elements in the even 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.

Groups- these 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 on the outer ns and np sublevels. Side subgroups consist of elements of large periods. Their valence electrons are located in 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, 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 increase, and non-metallic properties weaken. 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 group eight. Common to the elements of the main and secondary subgroups are the formulas of 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 properties. For elements of the main subgroups, the formulas for hydrogen compounds are common. Elements of groups I-III form solids- hydrides, since the oxidation state of hydrogen is -1. Elements of groups IV-VII are gaseous. Hydrogen compounds of 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.

Atomic radii, their periodic changes in the system of chemical elements

The radius of an atom decreases with increasing charges of atomic nuclei in a period, because the attraction of the electron shells by the nucleus increases. A kind of “compression” occurs. 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, and the sizes of atoms decrease. 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 located further from the nucleus are easily separated from it, which is typical for metal elements.

In the same group, as the period number increases, the atomic radii increase, because increasing the charge of an atom has the opposite effect. From the point of view of the theory of atomic structure, whether elements belong to metals or non-metals is determined by the ability of their atoms to give up or gain electrons. Metal atoms give up electrons relatively easily and cannot gain them to complete their outer electron layer.


D.I. Mendeleev formulated a periodic law in 1869, which sounds like this: the properties of chemical elements and the substances formed by them are periodically dependent on the relative atomic masses of the elements. Systematizing chemical elements based on their relative atomic masses, Mendeleev also paid great attention to the properties of the elements and the substances formed by them, 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 their charges atomic nuclei, And modern formulation periodic law is as follows: the properties of chemical elements and the substances formed by them are periodically dependent on the charges of their atomic nuclei. The periodicity in changes in the properties of 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 on the outer level that reflect the symbolism adopted in the periodic table.


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

  • When moving FROM RIGHT TO LEFT along PERIOD METAL properties of p-elements INCREASED. IN reverse direction- non-metallic ones increase. This is explained by the fact that to the right are elements whose electronic shells are closer to the octet. Elements on the right side of the period are less likely to give up their electrons to form metallic bonds and in chemical reactions in general.
  • For example, carbon is a more pronounced nonmetallic than its period neighbor boron, and nitrogen has even more pronounced nonmetallic properties than carbon. From left to right in a period, the nuclear charge also increases. Consequently, the attraction of valence electrons to the nucleus increases and their release becomes more difficult. On the contrary, the s-elements on the left side of the table have few electrons in the outer shell and a lower nuclear charge, which promotes the formation of a metallic bond. With the obvious exception of hydrogen and helium (their shells are close to complete or complete!), all s-elements are metals; p-elements can be both metals and non-metals, depending on whether they are on the left or right side of the table.
  • The d- and f-elements, as we know, have “reserve” electrons from the “penultimate” shells, which complicate the simple picture characteristic of the s- and p-elements. In general, d- and f-elements exhibit metallic properties much more readily.
  • The overwhelming number of elements are metals and only 22 elements are classified as non-metals: H, B, C, Si, N, P, As, O, S, Se, Te, as well as all halogens and inert gases. Some elements, due to the fact that they can exhibit only weak metallic properties, are classified as semimetals. What are semimetals? If you select p-elements from the Periodic Table and write them in a separate “block” (this is done in the “long” form of the table), you will find a pattern shown in The lower left part of the block contains typical metals, top right - typical nonmetals. Elements that occupy places on the boundary between metals and non-metals are called semimetals.
  • Semimetals are located approximately along the diagonal running through the p-elements from the upper left to the lower right corner of the Periodic Table
  • Semimetals have covalent crystal lattice in the presence of metallic conductivity (electrical conductivity). They either have insufficient valence electrons to form a full-fledged “octet” covalent bond(as in boron), or they are not held tightly enough (as in tellurium or polonium) due to the large size of the atom. Therefore, the bond in covalent crystals of these elements is partially metallic in nature. Some semimetals (silicon, germanium) are semiconductors. The semiconducting properties of these elements are explained by many complex reasons, but one of them is the significantly lower (although not zero) electrical conductivity, explained by the weak metallic bond. The role of semiconductors in electronic technology is extremely important.
  • When moving TOP DOWN along the groups METAL IS REINFORCED properties of elements. This is due to the fact that lower in the groups there are elements that already have quite a lot of filled electron shells. Their outer shells are further from the core. They are separated from the nucleus by a thicker “coat” of lower electron shells, and the electrons of the outer levels are held less tightly.

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

  • The reasons listed above explain why FROM LEFT TO RIGHT OXIDATING INCREASES properties, and when moving TOP TO DOWN - RESTORATIVE properties of elements.
  • The latter pattern applies even to such unusual elements as inert gases. From the “heavy” noble gases krypton and xenon, which are in the lower part of the group, it is possible to “select” electrons and form 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 oxidizing agent fluorine (F), and in the lower left corner is 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 elements, their ELECTRONEGATIVITY INCREASES Same FROM LEFT TO RIGHT, reaching a maximum for halogens. Not the least role in this is played by the degree of completeness of the valence shell, its proximity to the octet.
  • When moving TOP DOWN by groups ELECTRONEGATIVITY DECREASES. This is due to an increase in the number of electron shells, on the last of which electrons are attracted to the nucleus weaker and weaker.
  • c) Regularities associated with the sizes of atoms.
  • Atomic sizes (ATOMIC RADIUS) when moving FROM LEFT TO RIGHT along the period REDUCED. Electrons are increasingly attracted to the nucleus as the nuclear charge 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. On the contrary, the size of a fluorine atom is smaller than that of an oxygen atom.
  • When moving FROM TOP TO DOWN ATOMIC RADIUS elements GROWING, because more electron shells are filled.

d) Regularities associated with the valence of elements.

  • Elements of the same SUBGROUPS 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).
  • d-Elements exhibit many different valencies that cannot be accurately predicted by group number.
  • Not only elements, but also many of their compounds - oxides, hydrides, compounds with halogens - exhibit periodicity. For each GROUPS elements, you can write formulas for compounds that are periodically “repeated” (that is, can be written in the form of a generalized formula).

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

Changes in 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);
  • highest degree oxidation increases (equal to group number);
  • the number of electronic layers of atoms does not change;
  • metallic properties are reduced;
  • the non-metallic properties of the elements increases.

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

  • the charge of atomic nuclei 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 connection between the electrons of the outer layer and the nucleus decreases;
  • electronegativity decreases;
  • the metallicity of elements increases;
  • the non-metallicity of elements decreases.

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

Reference material for taking the test:

Mendeleev table

Solubility table

A change in the metallic properties of chemical elements will be similar to a change in their atomic radii. Therefore, in the main subgroups, metallic properties increase with increasing serial number, and in periods with increasing serial number, metallic properties decrease. Non-metallic properties, on the contrary, in the main subgroups decrease with increasing serial number, and in periods with increasing serial number they increase. in a number of elements of a certain period, the properties of basic oxides and their corresponding hydroxides are weakened, and acid properties in the same direction intensify. The transition from basic to acidic oxides, and, accordingly, from bases to acids, occurs in the period through amphoteric oxide or hydroxide. This pattern is valid for the second and third periods of the periodic system. For elements of long periods, complex patterns are observed. During the transition from one period to another, the configuration of the outer electronic layer is periodically repeated, with the same properties of chemical elements, their simple substances and their compounds. This is the main explanation of the meaning of D.I. Mendeleev’s periodic law.

How does the activity of a metal depend on its position in the periodic table and the values ​​of ionization potentials?

The values ​​of electrochemical potentials are a function of many variables and therefore exhibit a complex dependence on the position of metals in the periodic table. Thus, the oxidation potential of cations increases with an increase in the atomization energy of the metal, with an increase in the total ionization potential of its atoms, and with a decrease in the hydration energy of its cations.

Metals located at the beginning of periods are characterized by low values ​​of electrochemical potentials and occupy places on the left side of the voltage series. In this case, alternation (alkaline and alkaline earth metals reflects the phenomenon of diagonal similarity. Metals located closer to the middle of the periods are characterized by large potential values ​​and occupy places in the right half of the row. A consistent increase in the electrochemical potential (from −3.395 V for the Eu2+/Eu pair to +1.691 V for the Au+/Au pair) reflects a decrease in the reducing activity of metals (the ability to donate electrons) and an increase in the oxidizing ability of their cations (the ability to gain electrons). Thus, the strongest reducing agent is metallic europium, and the strongest oxidizing agent is gold cations Au+.

Hydrogen is traditionally included in the voltage series, since practical measurement of the electrochemical potentials of metals is carried out using a standard hydrogen electrode.

What are metallides?

Metallides - metal compounds, intermetallic phases, intermediate phases, chemical compounds metals among themselves. Connections adjoin the M. transition metals with non-metals (H, B, C, N, etc.). In such compounds there is a metallic bond. Metals are obtained by direct interaction of their components upon heating, through exchange decomposition reactions, etc. The formation of metals is observed when an excess component is isolated from solid solutions or as a result of ordering in the arrangement of atoms of the components of solid solutions.

What methods can be used to determine the composition of the formed metallide?

The composition of metals usually does not correspond to the formal valency of their components and can vary within significant limits. This is explained by the fact that ionic bonds are rare in metals, and metallic bonds predominate.

Zverev V. B-23

Independent work №7

What is corrosion?

Corrosion is the spontaneous destruction of metals as a result of chemical or physical-chemical interaction with environment. In general, this is the destruction of any material, be it metal or ceramics, wood or polymer. The cause of corrosion is the thermodynamic instability of structural materials to the effects of substances in the environment in contact with them. An example is oxygen corrosion of iron in water: 4Fe + 6H2O + 3O2 = 4Fe(OH)3. Hydrated iron oxide Fe(OH)3 is what is called rust.

in periods from left to right:

· the radius of atoms decreases;
· electronegativity of 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 electronic layers of atoms does not change;
· metallic properties decrease;
· non-metallic properties of elements are increased.

Changing some element characteristics in a group from top to bottom:
· the charge of atomic nuclei increases;
· the radius of 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 connection between the electrons of the outer layer and the nucleus decreases;
electronegativity decreases;
· metallicity of elements increases;
· non-metallicity of elements decreases.

Elements that are in the same subgroup are analogue elements, because they have some general properties(same higher valence, same forms of oxides and hydroxides, etc.). These general properties are explained by the structure of the outer electronic layer.

Read more about the patterns of changes in the properties of elements by periods and groups

The acid–base properties of hydroxides depend on which of the two bonds in the E–O–H chain is less strong.
If the E–O bond is less strong, then the hydroxide exhibits basic properties if O−H − acidic.
The weaker these bonds are, the greater the strength of the corresponding base or acid. The strength of the E–O and O–H bonds in the hydroxide depends on the distribution of electron density in the E–O–H chain. The latter is most strongly influenced by the oxidation state of the element and the ionic radius. An increase in the oxidation state of an element and a decrease in its ionic radius cause a shift in the electron density towards the atom
element in the chain E ← O ←N. This leads to a weakening of the O–H bond and strengthening of the E–O bond. Therefore, the basic properties of the hydroxide are weakened, and the acidic ones are enhanced.