The charges of the nuclei of elements in the periodic table continuously increase, and the properties of simple substances repeat periodically. How to explain this?
D.I. Mendeleev noticed that the properties of elements periodically repeat with increasing values of their mass numbers. He arranged the 63 elements discovered by that time in order of increasing atomic masses, taking into account their chemical and physical properties. Mendeleev believed that the periodic law he discovered was a reflection of deep patterns in the internal structure of matter; he stated the fact of periodic changes in the properties of elements, but did not know the reasons for the periodicity.
Further study of the structure of the atom showed that the properties of substances depend on the charge of the nucleus of the atoms, and elements can be systematized based on their electronic structure. The properties of simple substances and their compounds depend on the periodically repeating electronic configuration of the valence sublevel of the element's atoms. Therefore, “electronic analogues” are also “chemical analogues”.
Let us write down the electronic formulas of the atoms of the elements of the main subgroups of the second and seventh groups.
Elements of the second group have a general electronic formula of valence electrons ns 2. Let's write down their electronic formulas:
Be 1s 2 2s 2,
Mg 1s 2 2s 2 2p 6 3s 2,
Ca 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2,
Sr 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2.
Elements of the seventh group have a common electronic formula for valence electrons ns 2 np 5, and the complete electronic formulas are:
F 1s 2 2s 2 2p 5 ,
Cl 1s 2 2s 2 2p 6 3s 2 3p 5 ,
Br 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p5 ,
I 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p5 .
So, the electronic structures of atoms are periodically repeated for elements of the same group, therefore their properties are periodically repeated, since they depend mainly on the electronic configuration of the valence electrons. Elements of the same group have common properties, but there are also differences. This can be explained by the fact that although atoms have the same electronic structure of valence electrons, but these electrons are located at different distances from the nucleus, the force of attraction of them to the nucleus when moving from period to period weakens, the atomic radius increases, valence electrons become more mobile, which affects the properties of substances.
41. Based on the position of germanium, cesium and technetium in the periodic table, create formulas for the following compounds: meta and orthogermanic acids, cesium dihydrogen phosphate and technetium oxide, corresponding to its highest oxidation state. Draw the structural formulas of these compounds.
42. What is ionization energy? In what units is it expressed? How does the reduction activity of s- and p-elements in groups of the periodic table change with increasing atomic number? Why?
43. What is electronegativity? How does the electronegativity of elements change in the second and third periods in a group of the periodic system with increasing atomic number?
44. Based on the position of germanium, molybdenum and rhenium in the periodic table, compose the gross formulas of the following compounds: hydrogen compound of germanium, rhenium acid and molybdenum oxide corresponding to its highest oxidation state. Draw the structural formulas of these compounds.
45. What is electron affinity? In what units is it expressed? How does the oxidative activity of nonmetals change in a period and in a group of the periodic system with increasing atomic number? Motivate your answer with the atomic structure of the corresponding element.
46. Make up formulas for the oxides and hydroxides of elements of the third period of the periodic table, corresponding to their highest oxidation state. How does the chemical character of these compounds change when moving from sodium to chlorine?
47. Which of the elements of the fourth period - vanadium or arsenic - has more pronounced metallic properties? Which element forms a gaseous compound with hydrogen? Motivate your answer based on the structure of the atoms of these elements.
48. What elements form gaseous compounds with hydrogen? What groups of the periodic table are these elements in? Make up formulas for hydrogen and oxygen compounds of chlorine, tellurium and antimony, corresponding to their lowest and highest oxidation states.
49. Which element of the fourth period - chromium or selenium - has more pronounced metallic properties? Which of these elements forms a gaseous compound with hydrogen? Motivate your answer by the structure of chromium and selenium atoms.
50. What is the lowest oxidation state of chlorine, sulfur, nitrogen and carbon? Why? Make up formulas for aluminum compounds with these elements in their oxidation states. What are the names of the corresponding compounds?
51. Which of the p-elements of the fifth group of the periodic table - phosphorus or antimony - has more pronounced non-metallic properties? Which of the hydrogen compounds of these elements is the stronger reducing agent? Motivate your answer with the atomic structure of these elements.
52. Based on the position of the metal in the periodic table, give a motivated answer to the question; which of the two hydroxides is a stronger base: Ba(OH) 2 or Mg(OH) 2; Ca(OH) 2 or Fe(OH) 2; Cd(OH) 2 or Sr(OH) 2?
53. Why does manganese exhibit metallic properties, and chlorine – non-metallic? Motivate your answer by the electronic structure of the atoms of these elements. Write the formulas of oxides and hydroxides of chlorine and manganese.
54. What is the lowest oxidation state of hydrogen, fluorine, sulfur and nitrogen? Why? Make up formulas for calcium compounds with these elements in their oxidation states. What are the names of the corresponding compounds?
55. What are the lowest and highest oxidation states of silicon, arsenic, selenium and chlorine? Why? Make up formulas for compounds of these elements that correspond to these oxidation states.
56. To which family do elements belong, in the atoms of which the last electron goes to the 4f and 5f orbitals? How many elements does each of these families contain?
57. The atomic masses of elements in the periodic table are continuously increasing, while the properties of simple bodies change periodically. How can this be explained?
58. What is the modern formulation of the periodic law? Explain why in the periodic table of elements argon, cobalt, tellurium and thorium are placed respectively before potassium, nickel, iodine and protactinium, although they have a larger atomic mass?
59. What are the lowest and highest oxidation states of carbon, phosphorus, sulfur and iodine? Why? Make up formulas for compounds of these elements that correspond to these oxidation states.
The periodic system of chemical elements is a classification of chemical elements created by D. I. Mendeleev on the basis of the periodic law discovered by him in 1869.
D. I. Mendeleev
According to the modern formulation of this law, in a continuous series of elements arranged in order of increasing magnitude of the positive charge of the nuclei of their atoms, elements with similar properties periodically repeat.
The periodic table of chemical elements, presented in table form, consists of periods, series and groups.
At the beginning of each period (except for the first), the element has pronounced metallic properties (alkali metal).
Symbols for the color table: 1 - chemical sign of the element; 2 - name; 3 - atomic mass (atomic weight); 4 - serial number; 5 - distribution of electrons across layers.
As the atomic number of an element increases, equal to the positive charge of the nucleus of its atom, metallic properties gradually weaken and non-metallic properties increase. The penultimate element in each period is an element with pronounced non-metallic properties (), and the last is an inert gas. In period I there are 2 elements, in II and III - 8 elements, in IV and V - 18, in VI - 32 and in VII (not completed period) - 17 elements.
The first three periods are called small periods, each of them consists of one horizontal row; the rest - in large periods, each of which (except for the VII period) consists of two horizontal rows - even (upper) and odd (lower). Only metals are found in even rows of large periods. The properties of the elements in these series change slightly with increasing ordinal number. The properties of elements in odd rows of large periods change. In period VI, lanthanum is followed by 14 elements, very similar in chemical properties. These elements, called lanthanides, are listed separately below the main table. Actinides, the elements following actinium, are presented similarly in the table.
The table has nine vertical groups. The group number, with rare exceptions, is equal to the highest positive valency of the elements of this group. Each group, excluding the zero and eighth, is divided into subgroups. - main (located to the right) and secondary. In the main subgroups, as the atomic number increases, the metallic properties of the elements become stronger and the non-metallic properties weaken.
Thus, the chemical and a number of physical properties of elements are determined by the place that a given element occupies in the periodic table.
Biogenic elements, i.e. elements that are part of organisms and perform a certain biological role in it, occupy the top part of the periodic table. Cells occupied by elements that make up the bulk (more than 99%) of living matter are colored blue; cells occupied by microelements are colored pink (see).
The periodic table of chemical elements is the greatest achievement of modern natural science and a vivid expression of the most general dialectical laws of nature.
See also, Atomic weight.
The periodic system of chemical elements is a natural classification of chemical elements created by D. I. Mendeleev on the basis of the periodic law discovered by him in 1869.
In its original formulation, D.I. Mendeleev’s periodic law stated: the properties of chemical elements, as well as the forms and properties of their compounds, are periodically dependent on the atomic weights of the elements. Subsequently, with the development of the doctrine of the structure of the atom, it was shown that a more accurate characteristic of each element is not the atomic weight (see), but the value of the positive charge of the nucleus of the element’s atom, equal to the serial (atomic) number of this element in the periodic system of D. I. Mendeleev . The number of positive charges on the nucleus of an atom is equal to the number of electrons surrounding the nucleus of the atom, since atoms as a whole are electrically neutral. In the light of these data, the periodic law is formulated as follows: the properties of chemical elements, as well as the forms and properties of their compounds, are periodically dependent on the magnitude of the positive charge of the nuclei of their atoms. This means that in a continuous series of elements arranged in order of increasing positive charges of the nuclei of their atoms, elements with similar properties will periodically repeat.
The tabular form of the periodic table of chemical elements is presented in its modern form. It consists of periods, series and groups. A period represents a successive horizontal series of elements arranged in order of increasing positive charge of the nuclei of their atoms.
At the beginning of each period (except for the first) there is an element with pronounced metallic properties (alkali metal). Then, as the serial number increases, the metallic properties of the elements gradually weaken and the non-metallic properties increase. The penultimate element in each period is an element with pronounced non-metallic properties (halogen), and the last is an inert gas. The first period consists of two elements, the role of an alkali metal and a halogen here is simultaneously played by hydrogen. Periods II and III include 8 elements each, called typical by Mendeleev. Periods IV and V contain 18 elements each, VI-32. The VII period has not yet been completed and is replenished with artificially created elements; There are currently 17 elements in this period. Periods I, II and III are called small, each of them consists of one horizontal row, IV-VII are large: they (with the exception of VII) include two horizontal rows - even (upper) and odd (lower). In even rows of large periods there are only metals, and the change in the properties of elements in the row from left to right is weakly expressed.
In odd series of large periods, the properties of the elements in the series change in the same way as the properties of typical elements. In the even row of the VI period, after lanthanum, there are 14 elements [called lanthanides (see), lanthanides, rare earth elements], similar in chemical properties to lanthanum and to each other. A list of them is given separately below the table.
The elements following actinium - actinides (actinides) - are listed separately and listed below the table.
In the periodic table of chemical elements, nine groups are located vertically. The group number is equal to the highest positive valency (see) of the elements of this group. The exceptions are fluorine (can only be negatively monovalent) and bromine (cannot be heptavalent); in addition, copper, silver, gold can exhibit a valency greater than +1 (Cu-1 and 2, Ag and Au-1 and 3), and of the elements of group VIII, only osmium and ruthenium have a valence of +8. Each group, with the exception of the eighth and zero, is divided into two subgroups: the main one (located to the right) and the secondary one. The main subgroups include typical elements and elements of long periods, the secondary subgroups include only elements of long periods and, moreover, metals.
In terms of chemical properties, the elements of each subgroup of a given group differ significantly from each other, and only the highest positive valency is the same for all elements of a given group. In the main subgroups, from top to bottom, the metallic properties of elements are strengthened and non-metallic ones are weakened (for example, francium is the element with the most pronounced metallic properties, and fluorine is non-metallic). Thus, the place of an element in Mendeleev’s periodic system (ordinal number) determines its properties, which are the average of the properties of neighboring elements vertically and horizontally.
Some groups of elements have special names. Thus, the elements of the main subgroups of group I are called alkali metals, group II - alkaline earth metals, group VII - halogens, elements located behind uranium - transuranium. Elements that are part of organisms, take part in metabolic processes and have a clear biological role are called biogenic elements. All of them occupy the top part of D.I. Mendeleev’s table. These are primarily O, C, H, N, Ca, P, K, S, Na, Cl, Mg and Fe, which make up the bulk of living matter (more than 99%). The places occupied by these elements in the periodic table are colored light blue. Biogenic elements, of which there are very few in the body (from 10 -3 to 10 -14%), are called microelements (see). The cells of the periodic system, colored yellow, contain microelements, the vital importance of which for humans has been proven.
According to the theory of atomic structure (see Atom), the chemical properties of elements depend mainly on the number of electrons in the outer electron shell. The periodic change in the properties of elements with an increase in the positive charge of atomic nuclei is explained by the periodic repetition of the structure of the outer electron shell (energy level) of the atoms.
In small periods, with an increase in the positive charge of the nucleus, the number of electrons in the outer shell increases from 1 to 2 in period I and from 1 to 8 in periods II and III. Hence the change in the properties of elements in the period from an alkali metal to an inert gas. The outer electron shell, containing 8 electrons, is complete and energetically stable (elements of group zero are chemically inert).
In long periods in even rows, as the positive charge of the nuclei increases, the number of electrons in the outer shell remains constant (1 or 2) and the second outer shell is filled with electrons. Hence the slow change in the properties of elements in even rows. In the odd series of large periods, as the charge of the nuclei increases, the outer shell is filled with electrons (from 1 to 8) and the properties of the elements change in the same way as those of typical elements.
The number of electron shells in an atom is equal to the period number. Atoms of elements of the main subgroups have a number of electrons in their outer shells equal to the group number. Atoms of elements of side subgroups contain one or two electrons in their outer shells. This explains the difference in the properties of the elements of the main and secondary subgroups. The group number indicates the possible number of electrons that can participate in the formation of chemical (valence) bonds (see Molecule), therefore such electrons are called valence. For elements of side subgroups, not only the electrons of the outer shells are valence, but also those of the penultimate ones. The number and structure of electron shells are indicated in the accompanying periodic table of chemical elements.
The periodic law of D.I. Mendeleev and the system based on it are of exceptionally great importance in science and practice. The periodic law and system were the basis for the discovery of new chemical elements, the precise determination of their atomic weights, the development of the doctrine of the structure of atoms, the establishment of geochemical laws of distribution of elements in the earth's crust and the development of modern ideas about living matter, the composition of which and the patterns associated with it are in accordance with the periodic system. The biological activity of elements and their content in the body are also largely determined by the place they occupy in Mendeleev’s periodic table. Thus, with an increase in the serial number in a number of groups, the toxicity of elements increases and their content in the body decreases. The periodic law is a clear expression of the most general dialectical laws of the development of nature.
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The periodic repetition of the properties of elements with increasing atomic number becomes especially clear if the elements are arranged in a table called the periodic table or periodic table of elements. Several forms of the periodic table have been proposed and are used.
The periodic repetition of the properties of elements with increasing atomic number can be clearly shown by arranging the elements in a table called the periodic table, or periodic table, of the elements. Many different forms of the periodic table have been proposed and are in use.
The principle of periodic repetition of the properties of elements could not allow the existence of only one, isolated element of argon; There must be several or none of these simple substances. However, Ramsay firmly stood on the position of the periodic law, and this, as well as the development of laboratory technology at the end of the last century, predetermined the rapid discovery of the remaining members of the group of inert gases.
What explains the periodic repetition of the properties of elements in the periodic table.
What explains the periodic repetition of the properties of elements.
Accepting that the periodic repetition of the properties of elements is due not only to their mass (atomic weight), but also to the nature of the movement of the atoms themselves as whole particles (the speed and direction of their movement), Flavitsky builds his hypothesis on the following basis: the periodicity of elements is not explained by what is repeated type of internal Structure of atoms, but by the fact that the nature of the movement of atoms as whole particles periodically changes.
Thus, the reason for the periodic repetition of the properties of elements is the periodic repetition of the electronic configurations of their atoms.
The study of the electronic structure of atoms made it possible to prove that the reason for the periodic repetition of the properties of elements with increasing atomic number is the periodic repetition of the process of building new electronic shells. The same group of the periodic table always includes those elements whose atoms have the same number of electrons in their outer shells. Thus, the atoms of all inert gases, except helium, contain 8 electrons in the outer shell and are most difficult to ionize, while atoms of alkali metals contain one electron in the outer shell and have the lowest ionization potential. Alkali metals with only one electron in their outer shell can easily lose it, becoming a stable positive ion form with an electron configuration similar to the nearest noble gas with a lower atomic number. Elements such as fluorine, chlorine, etc., which in terms of the number of external electrons approach the configuration of inert gases, on the contrary, tend to acquire electrons and reproduce this electronic configuration, turning into the corresponding negative ion.
The periods following the third of D.I. Mendeleev’s table are longer. However, the periodic repetition of element properties is preserved. It acquires a more complex character, due to the increasing diversity of physical and chemical properties of elements as their atomic masses increase. Consideration of the structure of atoms of the first periods confirms that the limited number of places for electrons in each shell (Pauli exclusion) surrounding the nucleus is the reason for the periodic repetition of the properties of elements. This periodicity is the great law of nature, discovered by D.I. Mendeleev at the end of the last century, and in our time has become one of the foundations for the development of not only chemistry, but also physics.
The values of /j gradually increase with increasing Z until Z reaches the noble gas value and then drops to about one-fourth of the noble gas value as we move to the next element. The frequency of changes in another property - the density of elements in the solid state - is shown in Fig. 5.13. This periodic repetition of the properties of elements with increasing serial number becomes especially clear if the elements are arranged in the form of a table called the periodic table or periodic system of elements. Many different forms of the periodic table have been proposed and are in use.
At the same time as Newlands, de Chancourtois was approaching the discovery of the periodic law in France. But in contrast to the sensual musical and sound image, which served as an analogy for Newlands with the pattern of chemical elements that he partially identified, the French naturalist used an abstract geometric image: he compared the periodic repetition of the properties of elements arranged according to the magnitude of their atomic weights with the winding of a spiral line (vis tellurique) and the side surface of the cylinder.
The idea of the magnitude of the nuclear charge as a defining property of the atom formed the basis of the modern formulation of D.I. Mendeleev’s periodic law: the properties of chemical elements, as well as the forms and properties of compounds of these elements, are periodically dependent on the magnitude of the charge of the nuclei of their atoms. It made it possible to explain the reason for the periodic repetition of the properties of elements, which lies in the periodic repetition of the structure of the electronic configurations of atoms.
Only after the structure of the atom was clarified, the reasons for the periodic repetition of the properties of elements became clear.
Periodic table of chemical elements by D. I. Mendeleev
Basic concepts:
1. Serial number of a chemical element- the number given to the element when numbering it. Shows the total number of electrons in an atom and the number of protons in the nucleus, determines the charge of the nucleus of an atom of a given chemical element.
2. Period– chemical elements arranged in a row (only 7 periods). The period determines the number of energy levels in an atom.
Small periods (1 – 3) include only s - and p - elements (elements of the main subgroups) and consist of one line; large ones (4 – 7) include not only s - and p - elements (elements of the main subgroups), but also d - and f - elements (elements of the secondary subgroups) and consist of two lines.
3. Groups– chemical elements arranged in a column (there are only 8 groups). The group determines the number of external level electrons for elements of the main subgroups, as well as the number of valence electrons in an atom of a chemical element.
Main subgroup (A)– includes elements of large and small periods (only s - and p - elements).
Side subgroup (B)– includes elements of only large periods (only d - or f - elements).
4. Relative atomic mass (A r) – shows how many times a given atom is heavier than 1/12 of a 12 C atom; this is a dimensionless value (a rounded value is used for calculations).
5. Isotopes- a variety of atoms of the same chemical element, differing from each other only in their mass, with the same atomic number.
Atomic structure
Basic concepts:
1. Electronic cloud is a quantum mechanics model that describes the movement of an electron in an atom.
2. Orbital (s, p, d, f) – part of the atomic space in which the probability of finding a given electron is greatest (~ 90%).
3. Energy level– this is an energy layer with a certain energy level of the electrons located on it.
The number of energy levels in an atom of a chemical element is equal to the number of the period in which this element is located.
4. The maximum possible number of electrons at a given energy level is determined by the formula:
N = 2 n 2 , where n is the period number
5. The distribution of orbitals by level is represented by the diagram:
6. Chemical element- This is a type of atom with a certain nuclear charge.
7. Composition atom :
|
Particle |
Charge |
Weight |
||
|
Cl |
conventional units |
a.e.m. |
||
|
Electron (ē) |
1.6 ∙ 10 -19 |
9.10 ∙ 10 -28 |
0.00055 |
|
|
Proton ( p) |
1.6 ∙ 10 -19 |
1.67 ∙ 10 -24 |
1.00728 |
|
|
Neutron ( n) |
1.67 ∙ 10 -24 |
1.00866 |
||
8. Composition atomic nucleus:
The composition of the nucleus includes elementary particles -
protons(p) and neutrons(n).
· Because Almost all the mass of an atom is concentrated in the nucleus, then rounded valueA rof a chemical element is equal to the sum of protons and neutrons in the nucleus.
9. The total number of electrons in the electron shell of an atom is equal to the number of protons in the nucleus and the atomic number of the chemical element.
The order of filling levels and sublevels with electrons
I. The electronic formulas of atoms of chemical elements are in the following order:
· First, the total number of electrons in an atom is determined by the element number in D.I. Mendeleev’s table;
· Then, by the number of the period in which the element is located, the number of energy levels is determined;
· Levels are divided into sublevels and orbitals, and filled with electrons in accordance The principle of least energy
· For convenience, electrons can be distributed among energy levels using the formula N = 2n 2 and taking into account the fact that:
1. at the elements main subgroups(s -; p -elements) the number of electrons in the outer level is equal to the group number.
2. at the elements side subgroups usually at the external level two electron (with the exception of atoms Cu, Ag, Au, Cr, Nb, Mo, Ru, Rh, which at the external level one electron, y Pd at the external level zero electrons);
3. the number of electrons in the penultimate level is equal to the total number of electrons in the atom minus the number of electrons in all other levels.
II. The order in which electrons fill atomic orbitals is determined:
1.Principle of least energy
Energy scale:
1s<2s<2p<3s<3p<4s<3d<4p<5s<4d<5p<6s<4f<5d<6p<7s…
2. The state of an atom with a completely or half-filled sublevel (i.e., when each orbital has one unpaired electron) is more stable.
This explains the “failure” of the electron. Thus, the stable state of the chromium atom corresponds to the following electron distribution:
Cr: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5, not 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4,
i.e., an electron “fails” from the 4s sublevel to the 3d sublevel.
III. Families of chemical elements.
Elements in whose atoms the s-sublevel is filled with electrons external s-elements. These are the first 2 elements of each period, constituting the main subgroups I And II groups.
Elements in whose atoms the p-sublevel is filled with electrons external energy level are called p-elements. These are the last ones 6 elements of each period (except I And VII), making up the main subgroups III- VIII groups.
Elements in which the d-sublevel is filled second outside the level are called d-elements. These are elements of inserted decades IV, V, VI periods.
Elements in which the f-sublevel is filled third outside the level are called f-elements. The f elements include lanthanides and actinides.
Periodic law of D. I. Mendeleev
The properties of simple substances, as well as the forms and properties of compounds of elements, are periodically dependent on the atomic weights of the elements.
Modern formulation of the periodic law.
The properties of chemical elements and their compounds are periodically dependent on the magnitude of the charge of the nuclei of their atoms, expressed in the periodic repeatability of the structure of the outer valence electron shell.
Basic provisions
1. In the period from left to right:
2) Core charge – increases
3) Number of energy levels – constant
4) The number of electrons at the external level increases
5) Radius of atoms – decreases
6) Electronegativity – increases
Consequently, the outer electrons are held tighter, and the metallic (reducing) properties are weakened and the non-metallic (oxidizing) properties are enhanced.
2. In the group, in the main subgroup from top to bottom:
1) Relative atomic mass – increases
2) The number of electrons at the external level is constant
3) Core charge – increases
4) Number of energy levels – increases
5) Radius of atoms - increases
6) Electronegativity – decreases.
Consequently, the outer electrons are held weaker, and the metallic (reducing) properties of the elements are enhanced, while the non-metallic (oxidizing) properties are weakened.
3. Changes in the properties of volatile hydrogen compounds:
1) in groups of main subgroups, with increasing nuclear charge, the strength of volatile hydrogen compounds decreases, and the acidic properties of their aqueous solutions increase (the basic properties decrease);
2) in periods from left to right, the acidic properties of volatile hydrogen compounds in aqueous solutions increase (the basic ones decrease), and the strength decreases;
3) in groups with increasing nuclear charge in the main subgroups, the valence of the element in volatile hydrogen compounds does not change; in periods from left to right it decreases from IV to I.
4. Changes in the properties of higher oxides and their corresponding hydroxides (oxygen-containing acids of non-metals and metal bases):
1) in periods from left to right, the properties of higher oxides and their corresponding hydroxides change from basic through amphoteric to acidic;
2) the acidic properties of higher oxides and their corresponding hydroxides increase with increasing nuclear charge in the period, the basic properties decrease, and the strength decreases;
3) in the groups of main subgroups of higher oxides and their corresponding hydroxides, with increasing nuclear charge, the strength increases, the acidic properties decrease, and the basic properties increase;
4) in groups with increasing nuclear charge in the main subgroups, the valence of the element in higher oxides does not change; in periods from left to right it increases from I to VIII.
5. Completeness of the external level - if there are 8 electrons at the external level of the atom (for hydrogen and helium 2 electrons)
6. Metallic properties – the ability of an atom to donate electrons before completing the outer level.
7. Non-metallic properties - the ability of an atom to accept electrons before completing the outer level.
8. Electronegativity – the ability of an atom in a molecule to attract electrons to itself
9. Families of elements:
Alkali metals (1 group “A”) –Li, Na, K, Rb, Cs, Fr
Halogens (group 7 “A”) –F, Cl, Br, I
Inert gases (8th group “A”) –He, Ne, Ar, Xe, Rn
Chalcogens (group 6 “A”) –O, S, Se, Te, Po
Alkaline earth metals (group 2 “A”) –Ca, Sr, Ba, Ra
10. Atomic radius – distance from the atomic nucleus to the outer level
Tasks for consolidation:
Plan.
1. Periodic law D.I. Mendeleev and his general scientific and philosophical significance.
2. The periodic system and the serial number of an element as its most important characteristic. Periods and groups.
3.Changing the properties of elements in the periodic table.
4.The location of metals and non-metals in the periodic table.
1. Periodic law (D.I. Mendeleev, 1869)
The properties of elements, as well as the forms and properties of their compounds, periodically depend on the magnitude of the charge of the nuclei of their atoms
Why do the properties of elements repeat periodically?
With increasing nuclear charge in elements the number and distribution of valence electrons is periodically repeated, on which the properties of elements largely depend
2. Periodic table of elements
This is a graphical representation of the periodic law. In the periodic table, horizontal (period) and vertical (group) directions are distinguished.
Period
A horizontal row of elements in which the same number of energy levels are filled with electrons. III period: Na, Mg, Al, Si, P, S, Cl, Ar – the atoms of these elements fill 3 energy levels. There are 7 periods in the periodic system: 1,2,3 – small (consist of one row); 4,5,6,7 – large (have two rows); 7th period – unfinished.
Group
A vertical row of elements that have the same number of valence electrons, equal to the group number, and the same maximum valency. There are 8 groups in the system. Depending on how the valence electrons of the elements are distributed, the group is divided into two subgroups: main and secondary.
Subgroup
A vertical row of elements that have the same number and distribution of valence electrons, and therefore similar properties.
Main subgroup – group “A”
A vertical row of elements in which all valence electrons are located in the last level. The main subgroup includes elements of large and small periods.
Side subgroup "B"
A vertical row of elements in which, regardless of the group number, there are no more than 2 electrons at the last level, the remaining valence electrons are located at the penultimate level. The secondary subgroups include elements of only long periods
Periodic table and atomic structure
1. The atomic number of an element indicates the positive charge of the nucleus, the number of protons in the nucleus, and the number of electrons in the atom.
2. The period number indicates the number of energy levels in the atom.
3. Group numbers for all elements, with some exceptions, indicate the number of valence electrons, for elements of the main subgroups - the number of outer electrons.
3.
CHANGES IN THE PROPERTIES OF ELEMENTS IN THE PERIODIC SYSTEM
Atomic radius, r
In a period from left to right, the radius of the atom decreases slightly, because With the same number of energy levels, as a result of an increase in the charge of the nucleus, electrons are attracted more strongly. In the main subgroup from top to bottom, with an increase in the number of energy levels, the radius of the atom increases. In the side subgroup it changes nonlinearly.
Ionization energy, EI
This is the energy required to remove an electron from an atom. Expressed in electron volts. In the period with an increase in the charge of the nucleus, the number, external electrons, and a decrease in the radius of the atom from left to right, it increases; in the main subgroup, with an increase in the radius of the atom, it decreases from top to bottom.
Electron affinity energy, ES
The energy released when one electron is added to an atom. In the period from left to right it increases, in the main subgroup it decreases from top to bottom. Expressed in electron volts.
Electronegativity, EO
This is the ability of an atom in a molecule to attract electrons to itself. In the period from left to right it increases, in the main subgroup it decreases from top to bottom. Fluorine has the highest electronegativity value.
Number of electrons in the outer level
In a period, from left to right, it increases from I to 8 (the exception is the 1st period, from I to 2). Elements of the main subgroups are equal to the group number (with the exception of H, He), elements of the side subgroups have no more than 2 electrons at the outer level. When forming chemical compounds, atoms tend to a stable state - 8 electrons at the outer level (for the first elements - 2e). This is achieved by donating or adding electrons, depending on what is easier for the atom to do.
4.
METALS AND NON-METALS
IN THE PERIODIC CHART
Metals
Elements whose atoms at the outer energy level contain a small number of electrons: 1, 2, 3. When forming compounds, metals always give up ē and have only a positive charge.
Nonmetals
Elements whose atoms contain 4-8 electrons at the outer energy level. When forming compounds, nonmetals can either accept electrons (a negative charge arises) or give up electrons (a positive charge arises).
If in the periodic table we draw a diagonal from boron (Z = 5) to astatine (Z = 85), then down from the diagonal all elements are metals, and up are non-metals, with the exception of elements of side subgroups. Elements of side subgroups at the external level have no more than 2 ē, they all belong to metals.
There is no clear boundary between metals and non-metals; it is more correct to talk about metallicity and non-metallicity of an element.
Metallicity
The ability of an atom to give up electrons. In the period from left to right with increasing number ē and at the external level the metallicity weakens. In the main subgroups, metallicity increases from top to bottom, because The radius of the atom increases, the strength of the connection between external ē and the nucleus decreases, and the ability to give away ē increases.
Non-metallicity
The ability of an atom to gain electrons.
In the period from left to right with increasing number e at the external level increases; in the main subgroup, from top to bottom, it weakens with increasing atomic radius.
Thus, each period, with the exception of the first, begins with an active metal (alkali), ends with an active non-metal (halogen) and an inert gas. The most active metal is francium, the most active non-metal is fluorine.
