How oxidizing properties are enhanced in the periodic table. How the properties of chemical elements change in subgroups of the periodic system of Mendeleev

(Z) is periodic. Within the same period with increasing Z there is a tendency to a decrease in the size of atoms. For example, in the second period, atomic radii have the following values:

r , nm	0,155	0,113	0,091	0,077	0,071	0,066	0,064

This is explained by an increase in the attraction of the electrons of the outer layer to the nucleus as the charge of the nucleus increases. In subgroups, from top to bottom, atomic radii increase, because the number of electron layers increases:

	r , nm		r , nm
	0,155		0,071
	0,189		0,130
	0,236		0,148
	0,248		0,161
	0,268		0,182

The loss of electrons by an atom leads to a decrease in its effective size, and the addition of excess electrons leads to an increase. Therefore, the radius of a positive ion (cation) is always less, and the radius of a negative ion (anion) is always greater than the radius of the corresponding electrically neutral atom. For example:

	r , nm		r , nm
	0,236	Cl 0	0,099
	0,133	Cl -	0,181

The radius of the ion is the more different from the radius of the atom, the greater the charge of the ion:

	cr 0	Cr2+	Cr3+
r , nm	0,127	0,083	0,064

Within one subgroup, the radii of ions of the same charge increase with increasing nuclear charge:

	r , nm		r , nm
Li+	0,068		0,133
Na+	0,098	Cl -	0,181
	0,133	Br -	0,196
Rb+	0,149		0,220

This regularity is explained by the increase in the number of electron layers and the growing distance of the outer electrons from the nucleus.

b) Ionization energy and electron affinity. In chemical reactions, the nuclei of atoms do not undergo changes, while the electron shell is rebuilt, and the atoms are able to turn into positively and negatively charged ions. This ability can be quantified by the ionization energy of an atom and its electron affinity.

Ionization energy (ionization potential) I is the amount of energy required to detach an electron from an unexcited atom to form a cation:

X- e→ X+

Energy ionization is measured in kJ/mol or in electronvolts 1 eV = 1.602. 10 -19 J or 96.485 kJ / mol.(eV). The detachment of the second electron is more difficult than the first, because the second electron is detached not from a neutral atom, but from a positive ion:

X+- e→ X 2+

Therefore, the second ionization potential I 2 more than the first ( I 2 >I one). Obviously, the removal of each next electron will require more energy than the removal of the previous one. To characterize the properties of elements, the energy of detachment of the first electron is usually taken into account.

In groups, the ionization potential decreases with increasing atomic number of the element:

I, eV	6,39	5,14	4,34	4,18	3,89

This is due to the greater distance of valence electrons from the nucleus and, consequently, their easier detachment as the number of electron layers increases. The value of the ionization potential can serve as a measure of the “metallicity” of an element: the lower the ionization potential, the easier it is to remove an electron from an atom, the more pronounced the metallic properties.

In periods from left to right, the charge of the nucleus increases, and the radius of the atom decreases. Therefore, the ionization potential gradually increases, and the metallic properties weaken:

I, eV	5,39	9,32	8,30	11,26	14,53	13,61	17,42	21,56

Breaking the upward trend I observed for atoms with a completely filled external energy sublevel, or for atoms in which the external energy sublevel is exactly half filled:

This indicates an increased energy stability of electronic configurations with completely or exactly half-occupied sublevels.

The degree of attraction of an electron to the nucleus and, consequently, the ionization potential depends on a number of factors, and above all on nuclear charge The charge of the nucleus is equal to the ordinal number of the element in the periodic table., on the distance between the electron and the nucleus, on the screening effect of other electrons. So, for all atoms, except for the elements of the first period, the influence of the nucleus on the electrons of the outer layer is screened by the electrons of the inner layers.

The field of the nucleus of an atom, which holds the electrons, also attracts a free electron if it is near the atom. True, this electron experiences repulsion from the electrons of the atom. For many atoms, the energy of attraction of an additional electron to the nucleus exceeds the energy of its repulsion from the electron shells. These atoms can add an electron, forming a stable singly charged anion. The energy of detachment of an electron from a negative singly charged ion in the process X - - e → X 0 is called the affinity of an atom for an electron ( A), measured in kJ/mol or eV. When two or more electrons are attached to an atom, repulsion prevails over attraction - the affinity of an atom for two or more electrons is always negative. Therefore, monatomic multiply charged negative ions (O 2-, S 2-, N 3-, etc.) cannot exist in the free state.

Electron affinity is not known for all atoms. Halogen atoms have the highest electron affinity.

B) electronegativity. This value characterizes the ability of an atom in a molecule to attract binding electrons to itself. Electronegativity should not be confused with electron affinity: the first concept refers to an atom in a molecule, and the second to an isolated atom. Absolute electronegativity(kJ/mol or eV 1 electronvolt = 1.602. 10 -19 J or 96.485 kJ / mol.) is equal to the sum of the ionization energy and electron affinity :AEO= I+A. In practice, the relative value is often used electronegativity, equal to the ratio of the AEO of this element to the AEO of lithium (535 kJ/mol):

A.I. Khlebnikov, I.N. Arzhanova, O.A. Napilkova

The atomic radii of elements and ions are calculated on the basis of internuclear distances, which depend not only on the nature of the atoms, but also on the nature of the chemical bond between them and on the state of aggregation of the substance.

Radii of atoms and equally charged ions in a period with increasing charges, the nuclei mainly (with a few exceptions) decrease due to an increase in the Coulomb attraction forces due to an increase in the number, and, consequently, the total charge of electrons in electron shells and nuclei.

In subgroups, with increasing nuclear charge (moving from top to bottom), atomic and ionic radii, as a rule, increase, which is associated with an increase in the number of electronic levels.

Ionization energy (I) (ionization potential) in the period increases with the growth of the nuclear charge, in the main and third secondary subgroups it decreases from top to bottom due to the appearance of a new energy level. In the remaining side subgroups, the ionization energy increases with increasing nuclear charge.

Electron affinity (E) ( energy released when an additional electron is attached to an atom, ion or molecule). The maximum at halogen atoms. Electron affinity depends not only on the charge of the atomic nucleus, but also on the degree of filling of the outer electronic levels.

Electronegativity (EO)- a generalized characteristic of an element, defined as the sum of the ionization energy and electron affinity.

Relative EC according to Pauling is defined as the ratio of the EO of an element to the EO of a lithium atom. The relative electronegativity in the period increases, and in subgroups it decreases with increasing nuclear charge.

The oxidizing power of the element changes in the same way as the electronegativity, and the reducing power in the reverse order.

Density of simple substances in a period usually passes through a maximum lying approximately in the middle of the period, increases in subgroups with increasing nuclear charge.

Basic properties of higher oxides and hydroxides of elements in the period they naturally weaken, which is associated with an increase in the force of attraction of hydroxide ions to the central atom with an increase in the charge of its nucleus and a decrease in the atomic radius, and in a subgroup, as a rule, they increase, because the atomic radius of the elements increases.

Acid properties these compounds change in the opposite direction.

Non-metallic properties in the period, as a rule, they increase from left to right, and in the subgroup they weaken from top to bottom, metal - vice versa. The boundary between metals and non-metals in the table runs along the diagonal B-At in such a way that all non-metals are in the upper right part of the table (the exception is d-elements).

The mechanism of chemical bond formation can be modeled in various ways.

An increase in the oxidation state of an element and a decrease in the radius of its ion (in this case, a decrease in the effective negative charge on this oxygen) make the oxide more acidic. This explains the regular change in the properties of oxides from basic to amphoteric and further to acidic.

1) In one period, with an increase in the serial number, the acidic properties of oxides increase and the strength of their corresponding acids increases.

2) In the main subgroups of the periodic system, when moving from one element to another from top to bottom, an increase in the new properties of oxides is observed:

3) With an increase in the degree of oxidation of the element, the acidic properties of the oxide increase and the basic ones weaken.

Chemical properties of oxides

Basic oxides

The main oxides are:

Oxides of all metals of the main subgroup of the first group (alkali metals Li - Fr)

The main subgroup of the second group, starting with magnesium (Mg - Ra)

Transition metal oxides in lower oxidation states, eg MnO, FeO.

Most of the basic oxides are solid crystalline substances of an ionic nature; metal ions are located at the nodes of the crystal lattice, which are quite strongly associated with O2- oxide ions; therefore, oxides of typical metals have high melting and boiling points.

We note one characteristic feature of oxides. The proximity of the ionic radii of many metal ions leads to the fact that in the crystal lattice of oxides, part of the ions of one metal can be replaced by ions of another metal. This leads to the fact that the law of composition constancy often does not hold for oxides, and mixed oxides of variable composition may exist.

Most basic oxides do not decompose when heated, with the exception of oxides of mercury and noble metals:

2HgO \u003d 2Hg + O 2

2Ag2O = 4Ag + O2

When heated, basic oxides can react with acidic and amphoteric oxides, with acids:

BaO + SiO 2 \u003d BaSiO 3,

MgO + Al 2 O 3 \u003d Mg (AlO 2) 2,

ZnO + H 2 SO 4 \u003d ZnSO 4 + H 2 O.

Oxides of alkali and alkaline earth metals directly react with water:

Like other types of oxides, basic oxides can enter into redox reactions:

Fe 2 O 3 + 2Al \u003d Al 2 O 3 + 2Fe

3CuO + 2NH 3 \u003d 3Cu + N 2 + 3H 2 O

4FeO + O 2 \u003d 2Fe 2 + O 3

The main oxides of the most active metals (alkaline and alkaline earth, starting with calcium oxide) when interacting with water (hydration reaction) form their corresponding hydroxides (bases). For example, when calcium oxide (quicklime) dissolves in water, calcium hydroxide is formed - a strong base:

CaO + H 2 O → Ca (OH) 2

Basic oxides react with acids to form the corresponding salts:

CaO + 2HCl → CaCl 2 + H 2 O

The reaction of basic oxides with acidic oxides also leads to the formation of salts:

Na 2 O + CO 2 → Na 2 CO 3

And with amphoteric oxides:

Li 2 O + Al 2 O 3 → 2LiAlO 2

Acid oxides

Most non-metal oxides are acid oxides (CO2, SO3, P4O10). Transition metal oxides in higher oxidation states also predominantly exhibit the properties of acidic oxides, for example: CrO3, Mn2O7, V2O5.

Acid oxides are its non-metal or transition metal oxides in high oxidation states and can be obtained by methods similar to those for basic oxides, for example:

4P + 5O 2 \u003d 2P 2 + O 5

2ZnS + 3O 2 \u003d 2ZnO + 2SO 2

K 2 Cr 2 O 7 + H 2 SO 4 \u003d 2CrO 3 ↓ + K 2 SO 4 + H 2 O

Na 2 SiO 3 + 2HCl = 2NaCl + SiO 2 ↓ + H 2 O

Most acidic oxides react directly with water to form acids:

The most typical for acid oxides are their reactions with basic and amphoteric oxides, with alkalis:

P 2 O 5 + Al 2 O 3 \u003d 2AlPO 4

Ca (OH) 2 + CO 2 \u003d CaCO 3 ↓ + H 2 O.

It was mentioned above that acidic oxides can enter into numerous redox reactions, for example:

2SO 2 +O 2 2SO 3

SO 2 + 2H 2 S \u003d 3S + 2H 2 O,

4CrO 3 + C 2 H 5 OH \u003d 2Cr 2 O 3 + 2CO 2 + ZN 2 O

Almost all acid oxides, when interacting with water (hydration), form their corresponding acid hydroxides (oxygen-containing acids). For example, when sulfur oxide (VI) is dissolved in water, sulfuric acid is formed:

SO 3 + H 2 O → H 2 SO 4

Acid oxides can be obtained from the corresponding acid:

H 2 SiO 3 → SiO 2 + H 2 O

Amphoteric oxides

Amphotericity (from the Greek Amphoteros - both) - the ability of chemical compounds (oxides, hydroxides, amino acids) to exhibit both acidic properties and basic properties, depending on the properties of the second reagent involved in the reaction.

Amphoteric oxides react with strong acids to form salts of these acids. Such reactions are a manifestation of the main properties of amphoteric oxides, for example:

ZnO + H 2 SO 4 → ZnSO 4 + H 2 O

They also react with strong alkalis, thereby showing their acidic properties, for example:

ZnO + 2NaOH → Na 2 ZnO 2 + H 2 O

Amphoteric oxides can react with alkalis in two ways: in solution and in melt.

When reacted with an alkali in the melt, an ordinary medium salt is formed (as shown in the example above).

When reacting with alkali in solution, a complex salt is formed.

Al 2 O 3 + 2NaOH + 3H 2 O → 2Na (In this case, sodium tetrahydroxoalluminate is formed)

Each amphoteric metal has its own coordination number.

For Be and Zn, this is 4; for and Al is 4 or 6; for and Cr is 6 or (very rarely) 4;

Amphoteric oxides usually do not dissolve in water and do not react with it.

Amphoteric oxides have a dual nature: they are simultaneously capable of reactions involving both basic and acidic oxides, i.e. react with both acids and alkalis:

Al 2 O 3 + 6HCl \u003d 2AlCl 3 + ZH 2 O,

Al 2 O 3 + 2NaOH + ZH 2 O \u003d 2Na [Al (OH) 4].

Amphoteric oxides include aluminum oxide Al2O3, chromium (III) oxide Cr2O3, beryllium oxide BeO, zinc oxide ZnO, iron oxide (III) Fe2O3 and a number of others.

The ideally amphoteric oxide is H2O water, which dissociates to form equal amounts of hydrogen ions (acid properties) and hydroxide ions (basic properties). The amphoteric properties of water are clearly manifested during the hydrolysis of salts dissolved in it:

Cu 2+ + H 2 O Cu (OH) + + H +,

CO 3 2- + H 2 O HCO 3- + OH -.

Dmitry Ivanovich Mendeleev discovered the periodic law, according to which the properties of the elements and the elements they form change periodically. This discovery was graphically displayed in the periodic table. The table shows very well and clearly how the properties of the elements change over the period, after which they are repeated in the next period.

To solve task No. 2 of the Unified State Exam in chemistry, we just need to understand and remember which properties of the elements change in which directions and how.

All this is shown in the figure below.

From left to right, electronegativity, non-metallic properties, higher oxidation states, etc. increase. And the metallic properties and radii decrease.

From top to bottom, vice versa: the metallic properties and radii of atoms increase, while the electronegativity decreases. The highest oxidation state, corresponding to the number of electrons in the outer energy level, does not change in this direction.

Let's look at examples.

Example 1 In the series of elements Na→Mg→Al→Si
A) the radii of atoms decrease;
B) the number of protons in the nuclei of atoms decreases;
C) the number of electron layers in atoms increases;
D) the highest degree of oxidation of atoms decreases;

If we look at the periodic table, we will see that all the elements of a given series are in the same period and are listed in the order in which they appear in the table from left to right. To answer this kind of question, you just need to know a few patterns of changes in properties in the periodic table. So from left to right along the period, metallic properties decrease, non-metallic ones increase, electronegativity increases, ionization energy increases, and the radius of atoms decreases. From top to bottom, metallic and reducing properties increase in a group, electronegativity decreases, ionization energy decreases, and the radius of atoms increases.

If you were attentive, you already understood that in this case the atomic radii decrease. Answer A.

Example 2 In order of increasing oxidizing properties, the elements are arranged in the following order:
A. F→O→N
B. I→Br→Cl
B. Cl→S→P
D. F→Cl→Br

As you know, in Mendeleev's periodic table, oxidizing properties increase from left to right in a period and from bottom to top in a group. Option B just shows the elements of one group in order from bottom to top. So B fits.

Example 3 The valence of elements in the higher oxide increases in the series:
A. Cl→Br→I
B. Cs→K→Li
B. Cl→S→P
D. Al→C→N

In higher oxides, the elements show their highest oxidation state, which will coincide with the valency. And the highest degree of oxidation grows from left to right in the table. We look: in the first and second versions, we are given elements that are in the same groups, where the highest degree of oxidation and, accordingly, the valence in oxides does not change. Cl → S → P - are located from right to left, that is, on the contrary, their valence in the higher oxide will fall. But in the row Al→C→N, the elements are located from left to right, the valence in the higher oxide increases in them. Answer: G

Example 4 In the series of elements S→Se→Te
A) the acidity of hydrogen compounds increases;
B) the highest degree of oxidation of elements increases;
C) the valence of elements in hydrogen compounds increases;
D) the number of electrons in the outer level decreases;

Immediately look at the location of these elements in the periodic table. Sulfur, selenium and tellurium are in the same group, one subgroup. Listed in order from top to bottom. Look again at the diagram above. From top to bottom in the periodic table, metallic properties increase, radii increase, electronegativity, ionization energy and non-metallic properties decrease, the number of electrons at the outer level does not change. Option D is ruled out immediately. If the number of external electrons does not change, then the valence possibilities and the highest oxidation state also do not change, B and C are excluded.

Option A remains. We check for order. According to the Kossel scheme, the strength of oxygen-free acids increases with a decrease in the oxidation state of an element and an increase in the radius of its ion. The oxidation state of all three elements is the same in hydrogen compounds, but the radius grows from top to bottom, which means that the strength of acids also grows.
The answer is A.

Example 5 In order of weakening of the main properties, the oxides are arranged in the following order:
A. Na 2 O → K 2 O → Rb 2 O
B. Na 2 O → MgO → Al 2 O 3
B. BeO→BaO→CaO
G. SO 3 → P 2 O 5 → SiO 2

The main properties of oxides weaken synchronously with the weakening of the metallic properties of the elements forming them. And Me-properties weaken from left to right or from bottom to top. Na, Mg and Al are just arranged from left to right. Answer B.

One of the most important laws of nature is the periodic law, discovered in 1869 by Mendeleev, which he formulated as follows: "The properties of simple substances, as well as the forms and properties of compounds, are in a periodic dependence on the atomic weights of the elements."

With the development of quantum chemistry, the periodic law received a strict theoretical justification, and with it a new formulation: "The properties of simple substances, as well as the forms and properties of compounds of elements, are in a periodic dependence on the magnitude of the charges of the nuclei of their atoms."

Before Mendeleev, many tried to systematize the elements, Mayer (Germany) came closest. In 1864, in his book, he gave a table in which the elements were also arranged in ascending order of their atomic masses, but Mayer placed only 27 elements in this table, less than half known at that time. The merit of Mendeleev is that in his table there was a place not only for all known elements, but empty spaces were left for elements not yet discovered (ekabor - Sc, ekaaluminum - Ga, ekasilicon - Ge).

From the point of view of the electronic structure of the atom:

Period name the horizontal sequence of elements starting with the alkali metal and ending with the noble gas with the same maximum value of the main quantum number, equal to the number of the period.

The number of elements in a period is determined by the capacity of the sublevels.

group elements is a vertical collection of elements with the same electronic configuration and a certain chemical similarity. The group number (with the exception of I, II, VIII side subgroups) is equal to the sum of valence electrons.

In addition to division by periods (determined by the principal quantum number), there is a division into families, determined by the orbital quantum number. If an s-sublevel is filled in an element, then the s-family or s-element; p-sublevel - p‑element; d-sublevel - d-element; f-sublevel - f-element.

In the short period form of the periodic system, there are 8 groups, each of which is divided into the main and secondary subgroups. I and II main subgroups are filled with s-elements; III-VIII main subgroups - p-elements. d-elements are in side subgroups. f-elements are placed in separate groups.

Thus, each element in the periodic system of elements occupies a strictly defined place, which is marked by a serial number and is associated with the structure of the electron shells of the atom.

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

Experimental studies have established the dependence of the chemical and physical properties of elements on their position in the periodic system.

Ionization energy called the energy that must be expended to detach and remove an electron from an atom, ion or molecule . It is expressed in J or eV (1eV=1.6.10 -19 J).

The ionization energy is measure of restorative capacity atom. The lower the value of the ionization energy, the higher the reduction ability of the atom. Atoms lose an electron and become positively charged ions.

electron affinity is the energy released when an electron is attached to an atom, molecule or radical.

The electron affinity energy of atoms naturally changes in accordance with the nature of the electronic structures of the atoms of the elements. Periods from left to right the electron affinity and oxidizing properties of the elements increase. In groups from top to bottom, electron affinity tends to decrease.

Halogens have the highest electron affinity, because by attaching one electron to a neutral atom, it acquires the complete electron configuration of a noble gas.

The characteristic of which of the atoms is easier to give or add an electron is called electronegativity which is half the sum of the ionization energy and electron affinity.

Electronegativity increases from left to right for elements of each period and decreases from top to bottom for elements of the same PS group.

Atomic and ionic radii

Atoms and ions do not have strictly defined boundaries due to the wave nature of electrons. Therefore, the conditional radii of atoms and ions, connected to each other by chemical bonds in crystals, are determined.

The radii of metal atoms in periods with an increase in the ordinal number of elements decrease, because with the same number of electron layers, the charge of the nucleus increases, and, consequently, the attraction of electrons by it.

Within each group of elements, as a rule, the radii of atoms increase from top to bottom., because the number of energy levels increases. The radii of the ions are also in a periodic dependence on the atomic number of the element.

Example. How do the sizes of atoms change within a period, when moving from one period to another, and within the same group? What elements have the minimum and maximum size of an atom?

Inside the period (from left to right), the sizes of atoms decrease, because the charge of the nucleus increases and the electrons are more strongly attracted to the nucleus. In the main subgroups, the sizes of atoms increase, because. the number of electron layers increases. In side subgroups, such changes are less noticeable due to d-compression, and when moving from period V to period VI, there is even a decrease in the size of atoms due to f-compression.

According to these rules, the minimum size of an atom is helium, and the maximum cesium. Francium has no long-lived isotopes (the natural isotope is radioactive, with a half-life of 21 minutes).

Metals and non-metals. The division of elements and simple substances into metals and non-metals is to a certain extent conditional.

In terms of physical properties, metals are characterized by high thermal and electrical conductivity, negative temperature coefficient of conductivity, specific metallic luster, ductility, ductility, etc.

By chemical properties, metals are characterized by the main properties of oxides and hydroxides and reducing properties.

Similar differences in the properties of simple substances are associated with the nature of the chemical bond during their formation. A metallic bond in metals is formed with a deficiency of valence electrons, and a covalent bond in non-metals with a sufficient number of them. Proceeding from this, it is possible to draw a vertical boundary between the elements of groups IIIA and IV. On the left - elements with a deficit of valence electrons, on the right - with an excess. This is the Zintl border.

Example. How do typical metals differ from nonmetals? Why and how do metallic properties change with an increase in the ordinal number of elements?

In the periodic table of elements, there are mainly metals, there are few non-metals (22 in total). Metals include all s-elements. This is due to the presence of a small number of valence electrons (1 or 2) in them, as a result of this deficiency of electrons, a metallic bond is formed.

All d - and f -elements are also metals. When chemical bonds are formed, s-electrons of the outer energy level and part or all of the d-electrons of the penultimate level act as valence electrons in atoms of d-elements, and d-electrons participate in the formation of chemical bonds only after all external s- electrons. In addition, the screening effect of the nuclear charge contributes to the ease of removal of s-electrons. It consists in reducing the impact on the electron of the positive charge of the nucleus due to the presence of other electrons between the electron in question and the nucleus (these are d - or f -electrons).

In p-elements, there is a competition between an increase in the number of valence electrons (non-metallic properties) and screening of the nuclear charge (metallic properties are enhanced). In this regard, the p-elements in the subgroup from top to bottom increase the stability of lower oxidation states.

From right to left, the non-metallic properties of atoms increase along the period, due to an increase in the charge of the atom's nucleus and the difficulty of recoiling electrons. In the subgroup from top to bottom, the metallic properties increase, since the bond between the outer electrons and the nucleus weakens.

The properties of compounds are divided into acid-base and redox. The Periodic Table of the Elements explains these patterns well. Let's consider this on the example of hydroxides.

If the element has a low oxidation state (+1 or +2), for example, Na-O-H, then the Na-O bond is less strong than O-H and the bond is broken along a weaker bond.

Na-O-H  Na + + Oh - . The compound has basic properties.

If the oxidation state of the element is large (from +5 to +7), then the element-oxygen bond is stronger than the O-H bond and the compound has acidic properties. In nitric acid, the oxidation state of nitrogen is large (+5).

 H + + NO 3 -

Compounds in the oxidation state +3 and +4 exhibit amphoteric properties, i.e. Depending on the reaction partner, they can exhibit both acidic and basic properties. But there are exceptions Zn +2, Be +2, Sn +2, Pb +2, Ge +2 have an oxidation state of +2, but are amphoteric compounds.

By period from right to left, the highest oxidation state, equal to the group number, increases, therefore non-metallic and acidic properties increase.

By subgroup top down increase metallic and basic properties, because the size of the atom increases and the bond with the neighboring atom is weakened .

Thus, the periodic system allows us to analyze the position of simple substances in connection with the peculiarities of their properties (metals, non-metals).

The periodic law of Mendeleev makes it possible to determine the properties of simple substances in chemical compounds. For the first time, the prediction of properties was carried out by Mendeleev himself. He calculated the properties of those elements that have not yet been discovered.