The periodic table 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, elements with similar properties are periodically repeated in a continuous series of elements arranged in ascending order of the positive charge of the nuclei of their atoms.
The periodic table of chemical elements, presented in the form of a table, consists of periods, rows and groups.
At the beginning of each period (except for the first) there is an element with pronounced metallic properties (alkali metal).
Legend for the color table: 1 - chemical sign of the element; 2 - name; 3 - atomic mass (atomic weight); 4 - serial number; 5 - distribution of electrons over layers.
As the ordinal number of the element, equal to the value of the positive charge of the nucleus of its atom, increases, metallic properties gradually weaken and nonmetallic properties increase. The penultimate element in each period is an element with pronounced non-metallic properties (), and the last is an inert gas. In the first period there are 2 elements, in II and III - 8 elements each, in IV and V - 18 each, 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 (excluding the VII period) consists of two horizontal rows - even (upper) and odd (lower). Only metals are in the even rows of large periods. The properties of the elements in these rows change little with increasing serial number. The properties of elements in odd rows of large periods change. In the VI period, lanthanum was followed by 14 elements, very similar in chemical properties. These elements, called lanthanides, are listed separately below the main table. Actinides - elements following actinium are presented in the table in a similar way.
There are nine vertical groups in the table. 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 subdivided into subgroups. - main (located to the right) and secondary. In the main subgroups, with an increase in the serial number, the metallic properties of the elements increase and the non-metallic properties of the elements 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, that is, elements that make up organisms and play a certain biological role in it, occupy the upper part of the periodic table. The cells occupied by the elements that make up the bulk (more than 99%) of living matter are colored blue, the cells occupied by microelements are colored pink (see).
The periodic table of chemical elements is the largest achievement of modern natural science and a vivid expression of the most general dialectical laws of nature.
See also, Atomic Weight.
The periodic table 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 the original formulation, the periodic law of D. I. Mendeleev stated: the properties of chemical elements, as well as the forms and properties of their compounds, are periodically dependent on the value of the atomic weights of the elements. Later, with the development of the theory 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 atom of the element, equal to the ordinal (atomic) number of this element in the periodic system of D.I.Mendeleev ... The number of positive charges in the nucleus of an atom is equal to the number of electrons surrounding the nucleus of an 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 value of the positive charge of the nuclei of their atoms. This means that in a continuous series of elements, arranged in the order of increasing positive charges of the nuclei of their atoms, elements with similar properties will be periodically repeated.
The tabular form of the periodic table of chemical elements is presented in its modern form. It consists of periods, rows and groups. A period is a sequential horizontal row of elements arranged in ascending order of the 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 gradually weaken and the non-metallic properties of the elements increase. The penultimate element in each period is an element with pronounced non-metallic properties (halogen), and the last is an inert gas. Period I consists of two elements, the role of an alkali metal and a halogen is simultaneously played by hydrogen. The II and III periods each include 8 elements, named by Mendeleev as typical. IV and V periods have 18 elements each, VI-32. The VII period is not yet completed and is being replenished with artificially created elements; there are currently 17 elements in this period. I, II and III periods 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, only metals are found, and the change in the properties of elements in a row from left to right is weak.
In odd series of large periods, the properties of 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 under the table.
The elements following actinium - actinides (actinides) are separately written out and listed under the table.
In the periodic table of chemical elements, there are nine groups along the vertical lines. The group number is equal to the highest positive valency (see) of the elements of this group. The exceptions are fluorine (it happens only negatively monovalent) and bromine (it is not heptavalent); in addition, copper, silver, gold can exhibit a valency of more 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 valency of +8. Each group, with the exception of the eighth and zero, is divided into two subgroups: the main (located to the right) and the secondary. The main subgroups include typical elements and elements of large periods, in secondary ones - only elements of large 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 valence is the same for all elements of a given group. In the main subgroups, from top to bottom, the metallic properties of the elements increase and the non-metallic ones weaken (for example, francium is an element with the most pronounced metallic properties, and fluorine is non-metallic). Thus, the place of an element in the periodic system of Mendeleev (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. So, 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 - transuranic. Elements that are part of organisms, take part in metabolic processes and have a pronounced biological role, are called biogenic elements. All of them occupy the upper part of DI Mendeleev's table. This is 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, which 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 trace elements, the vital importance of which for humans has been proven.
According to the theory of the structure of atoms (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 atoms.
In small periods, with an increase in the positive charge of the nucleus, the number of electrons on 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 alkali metal to inert gas. The outer electron shell, containing 8 electrons, is complete and energetically stable (the elements of the zero group are chemically inert).
In large periods in even rows, with an increase in the positive charge of the nuclei, the number of electrons on the outer shell remains constant (1 or 2) and the second shell outside the shell is filled with electrons. Hence the slow change in the properties of elements in even rows. In 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 for typical elements.
The number of electron shells in an atom is equal to the number of the period. The atoms of the elements of the main subgroups have on the outer shells the number of electrons equal to the group number. Atoms of elements of secondary subgroups contain one or two electrons on the 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, valence is not only the electrons of the outer shells, but also of the penultimate ones. The number and structure of the electron shells are indicated in the attached periodic table of chemical elements.
DI Mendeleev's periodic law and the system based on it are extremely important 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 theory of the structure of atoms, the establishment of geochemical laws of the distribution of elements in the earth's crust and the development of modern ideas about living matter, the composition of which and the related regularities 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 the periodic system of Mendeleev. So, 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 vivid expression of the most general dialectical laws of the development of nature.
DI Mendeleev came to the conclusion that their properties should be determined by some fundamental general characteristics. As such a fundamental characteristic for a chemical element, he chose the atomic mass of the element and briefly formulated the periodic law (1869):
The properties of elements, as well as the properties of the simple and complex bodies formed by them, are periodically dependent on the values \u200b\u200bof the atomic weights of the elements.
Mendeleev's merit lies in the fact that he understood the manifested dependence as an objective law of nature, which his predecessors could not do. DI Mendeleev believed that the composition of compounds, their chemical properties, boiling and melting points, crystal structure, and the like are in periodic dependence on atomic mass. A deep understanding of the essence of periodic dependence gave Mendeleev the opportunity to draw several important conclusions and assumptions.
Modern periodic table
First, of the 63 elements known at that time, Mendeleev changed the atomic masses of almost 20 elements (Be, In, La, Y, Ce, Th, U). Secondly, he predicted the existence of about 20 new elements and left a place for them in the periodic table. Three of them, namely ekabor, ekaaluminium and ecasilicon, have been described in sufficient detail and with surprising accuracy. This was triumphantly confirmed over the next fifteen years, when the elements Gallium (ekaaluminium), scandium (ekabor), and Germanium (ecasilicium) were discovered.
The periodic law is one of the fundamental laws of nature. Its influence on the development of the scientific worldview can only be compared with the law of conservation of mass and energy or quantum theory. Even in the time of D.I.Mendeleev, the periodic law became the basis of chemistry. Further discoveries of the structure and the phenomenon of isotopy showed that the main quantitative characteristic of an element is not the atomic mass, but the nuclear charge (Z). In 1913, Moseley and Rutherford introduced the concept of "ordinal number of an element", numbered all the symbols in the periodic system and showed that the basis of the classification of elements is the ordinal number of an element, equal to the charge of the nuclei of their atoms.
This statement is now known as Moseley's law.
Therefore, the modern definition of the periodic law is formulated as follows:
The properties of simple substances, as well as the forms and properties of compounds of elements, are periodically dependent on the value of the charge of their atomic nuclei (or on the ordinal number of the element in the periodic system).
The electronic structures of the atoms of elements clearly show that with an increase in the nuclear charge, a regular periodic repetition of electronic structures occurs, and hence the repetition of the properties of elements. This is reflected in the periodic table of the elements, for which several hundred variants have been proposed. Most often, two forms of tables are used - abbreviated and expanded, - containing all known elements and having free space for not yet opened.
Each element occupies a certain cell in the periodic table, which indicates the symbol and name of the element, its serial number, relative atomic mass, and for radioactive elements in square brackets the mass number of the most stable or available isotope is given. In modern tables, some other reference information is often given: density, boiling and melting points of simple substances, etc.
Periods
The main structural units of the periodic system are periods and groups - natural aggregates into which chemical elements are divided by electronic structures.
A period is a horizontal sequential row of elements in whose atoms electrons fill the same number of energy levels.
The period number coincides with the number of the outer quantum level. For example, the element calcium (4s 2) is in the fourth period, that is, its atom has four energy levels, and the valence electrons are in the outer, fourth level. The difference in the filling sequence of both the outer and the electron layers closer to the nucleus explains the reason for the different lengths of the periods.
In the atoms of the s- and p-elements, the building of the external level is taking place, in the d-elements - the second outside, and in the f-elements - the third outside the energy level.
Therefore, the difference in properties is most clearly manifested in the neighboring s- or p-elements. In d- and especially f-elements of the same period, the difference in properties is less significant.
As already mentioned, on the basis of the number of the energy sublevel built up by electrons, the elements are combined into electronic families. For example, in periods IV-VI there are families that contain ten d-elements: 3d-family (Sc-Zn), 4d-family (Y-Cd), 5d-family (La, Hf-Hg). In the sixth and seventh periods, fourteen elements each make up f-families: 4f-family (Ce-Lu), which is called lanthanoid, and 5f-family (Th-Lr) - actinoid. These families are listed below the periodic table.
The first three periods are called small, or typical periods, since the properties of the elements of these periods are the basis for the distribution of all other elements into eight groups. All other periods, including the seventh, incomplete, are called large periods.
All periods, except for the first, begin with alkaline (Li, Na, K, Rb, Cs, Fr) and end, with the exception of the seventh, incomplete, with inert elements (He, Ne, Ar, Kr, Xe, Rn). Alkali metals have the same external electronic configuration ns 1, where n - period number. Inert elements, except for helium (1s 2), have the same structure of the outer electron layer: ns 2 np 6, that is, electronic analogues.
The considered regularity makes it possible to come to the conclusion:
Periodic repetition of the same electronic configurations of the outer electron layer is the reason for the similarity of physical and chemical properties in analogous elements, since it is the outer electrons of atoms that mainly determine their properties.
In small typical periods, with an increase in the serial number, a gradual decrease in metallic and an increase in nonmetallic properties is observed, since the number of valence electrons at the external energy level increases. For example, atoms of all elements of the third period have three electron layers. The structure of the two inner layers is the same for all elements of the third period (1s 2 2s 2 2p 6), and the structure of the outer, third, layer is different. When passing from each previous element to each subsequent one, the charge of the atomic nucleus increases by one and, accordingly, the number of external electrons increases. As a result, their attraction to the nucleus increases, and the radius of the atom decreases. This leads to a weakening of the metallic properties and the growth of non-metallic ones.
The third period begins with the very active metal sodium (11 Na - 3s 1), followed by slightly less active magnesium (12 Mg - 3s 2). Both of these metals belong to the 3s family. The first p-element of the third period is aluminum (13 Al - 3s 2 3p 1), the metallic activity of which is less than that of magnesium, has amphoteric properties, that is, in chemical reactions it can behave like a non-metal. This is followed by non-metals silicon (14 Si - 3s 2 3p 2), phosphorus (15 P - 3s 2 3p 3), sulfur (16 S - 3s 2 3p 4), chlorine (17 Cl - 3s 2 3p 5). Their non-metallic properties are enhanced from Si to Cl, which is an active non-metal. The period ends with the inert element argon (18 Ar - 3s 2 3p 6).
Within one period, the properties of elements change gradually, and during the transition from the previous period to the next, a sharp change in properties is observed, since the building of a new energy level begins.
The gradual change in properties is characteristic not only for simple substances, but also for complex compounds, as shown in Table 1.
Table 1 - Some properties of elements of the third period and their compounds
| Electronic family | s-elements | p-elements | ||||||
|---|---|---|---|---|---|---|---|---|
| Element symbol | Na | Mg | Al | Si | P | S | Cl | Ar |
| Nuclear charge of an atom | +11 | +12 | +13 | +14 | +15 | +16 | +17 | +18 |
| External electronic configuration | 3s 1 | 3s 2 | 3s 2 3p 1 | 3s 2 3p 2 | 3s 2 3p 3 | 3s 2 3p 4 | 3s 2 3p 5 | 3s 2 3p 6 |
| Atomic radius, nm | 0,189 | 0,160 | 0,143 | 0,118 | 0,110 | 0,102 | 0,099 | 0,054 |
| Maximum valence | I | II | III | IV | V | VI | Vii | — |
| Higher oxides and their properties | Na 2 O | MgO | Al 2 O 3 | SiO 2 | P 2 O 5 | SO 3 | Cl 2 O 7 | — |
| Basic properties | Amphoteric properties | Acidic properties | — | |||||
| Hydrates of oxides (bases or acids) | NaOH | Mg (OH) 2 | Al (OH) 3 | H 2 SiO 3 | H 3 PO 4 | H 2 SO 4 | HClO 4 | — |
| Base | Weak base | Amphoteric hydroxide | Weak acid | Medium strength acid | Strong acid | Strong acid | — | |
| Compounds with hydrogen | NaH | MgH 2 | AlH 3 | SiH 4 | PH 3 | H 2 S | HCl | — |
| Salty solid substances | Gaseous substances | — | ||||||
In longer periods, the metallic properties weaken more slowly. This is due to the fact that, starting from the fourth period, ten transition d-elements appear, in which not the external, but the second outside d-sublevel is built up, and on the outer layer of d-elements there are one or two s-electrons, which determine to a certain extent the properties of these elements. Thus, for d-elements, the pattern becomes somewhat more complicated. For example, in the fifth period, the metallic properties gradually decrease from alkaline Rb, reaching the minimum strength for metals of the platinum family (Ru, Rh, Pd).
However, after the inactive Ag silver is placed cadmium Cd, in which an abrupt increase in metallic properties is observed. Further, with an increase in the serial number of an element, non-metallic properties appear and gradually increase up to the typical non-metal iodine. This period, like all the previous ones, ends with an inert gas. A periodic change in the properties of elements within large periods allows them to be divided into two rows, in which the second part of the period repeats the first.
Groups
Vertical columns of elements in the periodic table - groups consist of subgroups: main and secondary, they are sometimes denoted by the letters A and B, respectively.
The main subgroups include s- and p-elements, and the secondary subgroups - d- and f-elements of large periods.
The main subgroup is a set of elements that are placed vertically in the periodic table and have the same configuration of the outer electron layer in the atoms.
As follows from the above definition, the position of an element in the main subgroup is determined by the total number of electrons (s- and p-) of the external energy level, equal to the group number. For example, sulfur (S - 3s 2 3p 4 ), whose atom contains six electrons at the outer level, belongs to the main subgroup of the sixth group, argon (Ar - 3s 2 3p 6 ) - to the main subgroup of the eighth group, and strontium (Sr - 5s 2 ) - to the IIA-subgroup.
Elements of one subgroup are characterized by similar chemical properties. As an example, consider the elements of IА and VІІА subgroups (Table 2). With an increase in the nuclear charge, the number of electronic layers and the radius of the atom increases, but the number of electrons at the external energy level remains constant: for alkali metals (subgroup IA) - one, and for halogens (subgroup VIIA) - seven. Since it is the outer electrons that most significantly affect the chemical properties, it is clear that each of the considered groups of analogous elements has similar properties.
But within the same subgroup, along with the similarity of properties, some change is observed. So, the elements of subgroup ІА are all, except for H - active metals. But with an increase in the radius of the atom and the number of electronic layers screening the influence of the nucleus on the valence electrons, the metallic properties increase. Therefore, Fr is a more active metal than Cs, and Cs is more active than R in, etc. And in subgroup VIIA, for the same reason, the non-metallic properties of elements are weakened with an increase in the serial number. Therefore, F is a more active non-metal in comparison with Cl, and Cl is a more active non-metal in comparison with Br, etc.
Table 2 - Some characteristics of elements ІА and VІІА-subgroups
| period | Subgroup IA | Subgroup VIIA | ||||||
|---|---|---|---|---|---|---|---|---|
| Element symbol | Core charge | Atom radius, nm | Element symbol | Core charge | Atom radius, nm | External electronic configuration | ||
| II | Li | +3 | 0,155 | 2 s 1 | F | +9 | 0,064 | 2 s 2 2 p 5 |
| III | Na | +11 | 0,189 | 3 s 1 | Cl | +17 | 0,099 | 3 s 2 3 p 5 |
| IV | K | +19 | 0,236 | 4 s 1 | Br | 35 | 0,114 | 4 s 2 4 p 5 |
| V | Rb | +37 | 0,248 | 5 s 1 | I | +53 | 0,133 | 5 s 2 5 p 5 |
| VI | Cs | 55 | 0,268 | 6 s 1 | At | 85 | 0,140 | 6 s 2 6 p 5 |
| Vii | Fr | +87 | 0,280 | 7 s 1 | — | — | — | — |
A side subgroup is a set of elements that are placed vertically in the periodic table and have the same number of valence electrons due to the building up of the external s- and the second outside d-energy sublevels.
All elements of the secondary subgroups belong to the d-family. These elements are sometimes referred to as transition metals. In side subgroups, the properties change more slowly, since in the atoms of d-elements, the electrons build up the second from the outside energy level, and only one or two electrons are on the external level.
The position of the first five d-elements (subgroups IIIB-VIIB) of each period can be determined using the sum of external s-electrons and d-electrons of the second outside level. For example, from the electronic formula of scandium (Sc - 4s 2 3d 1 ) it can be seen that it is located in a side subgroup (since it is a d-element) of the third group (since the sum of the valence electrons is three), and manganese (Mn - 4s 2 3d 5 ) is located in the secondary subgroup of the seventh group.
The position of the last two elements of each period (subgroups IB and IIB) can be determined by the number of electrons at the outer level, since in the atoms of these elements the previous level is completely complete. For example, Ag (5s 1 5d 10) is placed in a secondary subgroup of the first group, Zn (4s 2 3d 10) - in a secondary subgroup of the second group.
The Fe-Co-Ni, Ru-Rh-Pd and Os-Ir-Pt triads are located in a secondary subgroup of the eighth group. These triads form two families: iron and platinoids. In addition to these families, the family of lanthanides (fourteen 4f-elements) and the family of actinides (fourteen 5f-elements) are separately distinguished. These families belong to a secondary subgroup of the third group.
An increase in the metallic properties of elements in subgroups from top to bottom, as well as a decrease in these properties within one period from left to right, cause the appearance of a diagonal pattern in the periodic system. So, Be is very similar to Al, B to Si, Ti to Nb. This is clearly manifested in the fact that in nature these elements form similar minerals. For example, in nature, Te always occurs with Nb, forming minerals - titanoniobates.
The graphic representation of the Periodic Law is the Periodic Table (table). The horizontal rows of the system are called periods, and the vertical columns are called groups.
There are 7 periods in the system (table), and the number of the period is equal to the number of electronic layers in the atom of the element, the number of the external (valence) energy level, and the value of the principal quantum number for the highest energy level. Each period (except the first) begins with an s-element - an active alkali metal and ends with an inert gas, in front of which is a p-element - an active non-metal (halogen). If we move along the period from left to right, then with an increase in the nuclear charge of atoms of chemical elements of small periods, the number of electrons at the external energy level will increase, as a result of which the properties of elements change - from typically metallic (since at the beginning of the period there is an active alkali metal), through amphoteric (the element exhibits the properties of both metals and non-metals) to non-metallic (active non-metal - halogen at the end of the period), i.e. metallic properties gradually weaken and non-metallic properties increase.
In large periods, with an increase in the nuclear charge, the filling of electrons is more complicated, which explains the more complex change in the properties of elements in comparison with elements of small periods. So, in even rows of large periods, with an increase in the nuclear charge, the number of electrons at the outer energy level remains constant and equal to 2 or 1. Therefore, while the filling with electrons of the next level after the outer (second outside) level, the properties of elements in even rows change slowly. When passing to odd series, with an increase in the value of the nuclear charge, the number of electrons at the external energy level (from 1 to 8) increases, the properties of elements change in the same way as in small periods.
DEFINITION
The vertical columns in the Periodic Table are groups of elements with a similar electronic structure and are chemical analogs. Groups are denoted by Roman numerals from I to VIII. Main (A) and secondary (B) subgroups are distinguished, the first of which contain s- and p-elements, the second - d - elements.
The subgroup number A indicates the number of electrons in the external energy level (the number of valence electrons). For elements of B-subgroups, there is no direct relationship between the group number and the number of electrons at the external energy level. In A-subgroups, the metallic properties of the elements increase, and the non-metallic ones decrease with an increase in the charge of the nucleus of the element's atom.
There is a relationship between the position of the elements in the Periodic Table and the structure of their atoms:
- the atoms of all elements of the same period have an equal number of energy levels, partially or completely filled with electrons;
- the atoms of all elements of A subgroups have an equal number of electrons at the external energy level.
Characterization plan of a chemical element based on its position in the Periodic Table
Usually, the characteristic of a chemical element based on its position in the Periodic Table is given according to the following plan:
- indicate the symbol of the chemical element, as well as its name;
- indicate the serial number, period and group number (type of subgroup) in which the element is located;
- indicate the nuclear charge, mass number, the number of electrons, protons and neutrons in the atom;
- record the electronic configuration and indicate the valence electrons;
- Sketch electronic-graphic formulas for valence electrons in the ground and excited (if possible) states;
- indicate the family of the element, as well as its type (metal or non-metal);
- compare the properties of a simple substance with the properties of simple substances formed by neighboring elements in a subgroup;
- compare the properties of a simple substance with the properties of simple substances formed by elements adjacent to the period;
- indicate the formulas of higher oxides and hydroxides with a brief description of their properties;
- indicate the values \u200b\u200bof the minimum and maximum oxidation states of a chemical element.
Characterization of a chemical element for example magnesium (Mg)
Consider the characteristics of a chemical element using the example of magnesium (Mg) according to the plan described above:
1. Mg - magnesium.
2. The serial number - 12. The element is in the 3rd period, in the II group, A (main) subgroup.
3. Z \u003d 12 (nuclear charge), M \u003d 24 (mass number), e \u003d 12 (number of electrons), p \u003d 12 (number of protons), n \u003d 24-12 \u003d 12 (number of neutrons).
4. 12 Mg 1s 2 2s 2 2p 6 3s 2 - electronic configuration, valence electrons 3s 2.
5. Basic condition
Excited state
6.s-element, metal.
7. The highest oxide - MgO - exhibits basic properties:
MgO + H 2 SO 4 \u003d MgSO 4 + H 2 O
MgO + N 2 O 5 \u003d Mg (NO 3) 2
Magnesium hydroxide is the base Mg (OH) 2, which exhibits all the typical properties of bases:
Mg (OH) 2 + H 2 SO 4 \u003d MgSO 4 + 2H 2 O
8. The oxidation state is "+2".
9. The metallic properties of magnesium are more pronounced than that of beryllium, but weaker than that of calcium.
10. The metallic properties of magnesium are less pronounced than that of sodium, but stronger than that of aluminum (neighboring elements of the 3rd period).
Examples of problem solving
EXAMPLE 1
| The task | Characterize the chemical element sulfur on the basis of its position in the D.I. Mendeleev |
| Decision | 1. S - sulfur. 2. The serial number - 16. The element is in the 3rd period, in the VI group, A (main) subgroup. 3. Z \u003d 16 (nuclear charge), M \u003d 32 (mass number), e \u003d 16 (number of electrons), p \u003d 16 (number of protons), n \u003d 32-16 \u003d 16 (number of neutrons). 4.16 S 1s 2 2s 2 2p 6 3s 2 3p 4 - electronic configuration, valence electrons 3s 2 3p 4. 5. Basic condition Excited state 6.p-element, non-metal. 7. Higher oxide - SO 3 - exhibits acidic properties: SO 3 + Na 2 O \u003d Na 2 SO 4 8. The hydroxide corresponding to the higher oxide - H 2 SO 4 exhibits acidic properties: H 2 SO 4 + 2NaOH \u003d Na 2 SO 4 + 2H 2 O 9. The minimum oxidation state is "-2", the maximum is "+6" 10. The non-metallic properties of sulfur are less pronounced than that of oxygen, but stronger than that of selenium. 11. The non-metallic properties of sulfur are more pronounced than that of phosphorus, but weaker than that of chlorine (neighboring elements in the 3rd period). |
EXAMPLE 2
| The task | Characterize the chemical element sodium based on its position in the D.I. Mendeleev |
| Decision | 1. Na - sodium. 2. The serial number - 11. The element is in the 3rd period, in the I group, in the A (main) subgroup. 3. Z \u003d 11 (nuclear charge), M \u003d 23 (mass number), e \u003d 11 (number of electrons), p \u003d 11 (number of protons), n \u003d 23-11 \u003d 12 (number of neutrons). 4.11 Na 1s 2 2s 2 2p 6 3s 1 - electronic configuration, valence electrons 3s 1. 5. Basic condition 6.s-element, metal. 7. The highest oxide - Na 2 O - exhibits basic properties: Na 2 O + SO 3 \u003d Na 2 SO 4 As sodium hydroxide, the base NaOH corresponds, which exhibits all the typical properties of bases: 2NaOH + H 2 SO 4 \u003d Na 2 SO 4 + 2H 2 O 8. The oxidation state is "+1". 9. The metallic properties of sodium are more pronounced than that of lithium, but weaker than that of potassium. 10. The metallic properties of sodium are more pronounced than that of magnesium (a neighboring element of the 3rd period). |
Atom composition.
An atom consists of atomic nucleus and electronic shell.
The nucleus of an atom consists of protons ( p +) and neutrons ( n 0). Most hydrogen atoms have a single proton nucleus.
Number of protons N(p +) is equal to the nuclear charge ( Z) and the ordinal number of the element in the natural series of elements (and in the periodic table of elements).
N(p +) = Z
The sum of the number of neutrons N(n 0), denoted simply by the letter N, and the number of protons Z called massive number and denoted by the letter AND.
A = Z + N
The electron shell of an atom consists of electrons moving around the nucleus ( e -).
Number of electrons N(e -) in the electron shell of a neutral atom is equal to the number of protons Z at its core.
The mass of a proton is approximately equal to the mass of a neutron and is 1840 times the mass of an electron, so the mass of an atom is practically equal to the mass of a nucleus.
The shape of the atom is spherical. The radius of the nucleus is about 100,000 times smaller than the radius of the atom.
Chemical element - the type of atoms (a set of atoms) with the same nuclear charge (with the same number of protons in the nucleus).
Isotope - a set of atoms of one element with the same number of neutrons in the nucleus (or the type of atoms with the same number of protons and the same number of neutrons in the nucleus).
Different isotopes differ from each other in the number of neutrons in the nuclei of their atoms.
The designation of a single atom or isotope: (E is the symbol of an element), for example:.
The structure of the electron shell of an atom
Atomic orbital - the state of an electron in an atom. Orbital symbol -. An electron cloud corresponds to each orbital.
The orbitals of real atoms in the ground (unexcited) state are of four types: s, p, d and f.
Electronic cloud - a part of space in which an electron can be found with a 90 (or more) percent probability.
Note: sometimes the concepts of "atomic orbital" and "electron cloud" are not distinguished, calling both the "atomic orbital".
The electron shell of the atom is layered. Electronic layer formed by electron clouds of the same size. Orbitals of one layer form electronic ("energy") level, their energies are the same for the hydrogen atom, but different for other atoms.
Orbitals of the same type are grouped into electronic (energy) sublevels:
s- sublevel (consists of one s-orbital), symbol -.
p- sublevel (consists of three p
d- sublevel (consists of five d-orbitals), symbol -.
f- sublevel (consists of seven f-orbitals), symbol -.
The energies of the orbitals of one sublevel are the same.
When designating sublevels, the number of the layer (electronic layer) is added to the symbol of the sublevel, for example: 2 s, 3p, 5d means s- sub-level of the second level, p-sublevel of the third level, d-sublevel of the fifth level.
The total number of sublevels in one level is equal to the level number n... The total number of orbitals at one level is n 2. Accordingly, the total number of clouds in one layer is also n 2 .
Designations: - free orbital (without electrons), - orbital with an unpaired electron, - orbital with an electron pair (with two electrons).
The order of filling the orbitals of an atom with electrons is determined by three laws of nature (the formulations are given in a simplified manner):
1. The principle of least energy - electrons fill the orbitals in the order of increasing orbital energy.
2. Pauli's principle - on one orbital there can be no more than two electrons.
3. Hund's rule - within the sublevel, electrons first fill free orbitals (one at a time), and only then form electron pairs.
The total number of electrons in the electronic level (or in the electron layer) is 2 n 2 .
The distribution of sublevels by energy is expressed as follows (in the order of increasing energy):
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p ...
This sequence is clearly expressed in an energy diagram:
The distribution of the electrons of an atom over levels, sublevels and orbitals (electronic configuration of an atom) can be depicted as an electronic formula, an energy diagram, or, simplified, as a diagram of electronic layers ("electronic circuit").
Examples of the electronic structure of atoms:
Valence electrons - the electrons of the atom, which can take part in the formation of chemical bonds. For any atom, these are all external electrons plus those pre-external electrons, the energy of which is greater than that of the external ones. For example: a Ca atom has external electrons - 4 s 2, they are also valence; the Fe atom has outer electrons - 4 s 2, but it has 3 d 6, therefore the iron atom has 8 valence electrons. The valence electronic formula of the calcium atom is 4 s 2, and the iron atom - 4 s 2 3d 6 .
Periodic table of chemical elements of D. I. Mendeleev
(natural system of chemical elements)
Periodic law of chemical elements (modern formulation): the properties of chemical elements, as well as simple and complex substances formed by them, are periodically dependent on the value of the charge from atomic nuclei.
Periodic system - graphic expression of the periodic law.
Natural range of chemical elements - a series of chemical elements, arranged according to the increasing number of protons in the nuclei of their atoms, or, which is the same, according to the increasing charges of the nuclei of these atoms. The ordinal number of an element in this row is equal to the number of protons in the nucleus of any atom of this element.
The table of chemical elements is constructed by "cutting" the natural series of chemical elements into periods (horizontal rows of the table) and groupings (vertical columns of the table) of elements with a similar electronic structure of atoms.
Depending on the method of combining elements into groups, the table may be long-period (elements with the same number and type of valence electrons are collected in groups) and short-period (elements with the same number of valence electrons are collected in groups).
The groups of the short-period table are divided into subgroups ( the main and collateral) that match the groups of the long-period table.
All atoms of elements of the same period have the same number of electronic layers equal to the period number.
The number of elements in periods: 2, 8, 8, 18, 18, 32, 32. Most of the elements of the eighth period are obtained artificially, the last elements of this period have not yet been synthesized. All periods, except the first, begin with an element that forms an alkali metal (Li, Na, K, etc.), and end with an element that forms a noble gas (He, Ne, Ar, Kr, etc.).
In the short-period table there are eight groups, each of which is divided into two subgroups (main and secondary), in the long-period table there are sixteen groups, which are numbered in Roman numerals with the letters A or B, for example: IA, IIIB, VIA, VIIB. Group IA of the long period table corresponds to the main subgroup of the first group of the short period table; group VIIB - a side subgroup of the seventh group: the rest are similar.
The characteristics of chemical elements change naturally in groups and periods.
In periods (with an increase in the serial number)
- the charge of the nucleus increases,
- the number of external electrons increases,
- the radius of atoms decreases,
- the strength of the bond of electrons with the nucleus (ionization energy) increases,
- electronegativity increases,
- the oxidizing properties of simple substances are enhanced ("nonmetallicity"),
- the reducing properties of simple substances ("metallicity") weaken,
- weakens the basic character of hydroxides and corresponding oxides,
- the acidic character of hydroxides and corresponding oxides increases.
In groups (with increasing serial number)
- the charge of the nucleus increases,
- the radius of atoms increases (only in A-groups),
- the bond strength of electrons with the nucleus decreases (ionization energy; only in A-groups),
- decreases electronegativity (only in A-groups),
- the oxidizing properties of simple substances weaken ("non-metallic"; only in A-groups),
- the reducing properties of simple substances are enhanced ("metallicity"; only in A-groups),
- the basic character of hydroxides and corresponding oxides increases (only in A-groups),
- the acidic nature of hydroxides and corresponding oxides weakens (only in A-groups),
- the stability of hydrogen compounds decreases (their reductive activity increases; only in A-groups).
Problems and tests on the topic "Topic 9." The structure of the atom. DI Mendeleev's Periodic Law and Periodic Table of Chemical Elements (PSKhE) "."
- Periodic law - Periodic law and the structure of atoms 8-9 grade
You should know: the laws of filling the orbitals with electrons (the principle of least energy, Pauli's principle, Hund's rule), the structure of the periodic table of elements.You must be able to: determine the composition of an atom by the position of an element in the periodic system, and, conversely, find an element in the periodic system, knowing its composition; to depict the structure diagram, the electronic configuration of an atom, ion, and, conversely, to determine the position of a chemical element in the PSCE according to the diagram and electronic configuration; to characterize the element and the substances formed by it according to its position in the PSCE; determine changes in the radius of atoms, the properties of chemical elements and the substances formed by them within one period and one main subgroup of the periodic system.
Example 1. Determine the number of orbitals at the third electronic level. What are these orbitals?
To determine the number of orbitals, we use the formula N orbitals \u003d n 2, where n - level number. N orbitals \u003d 3 2 \u003d 9. One 3 s-, three 3 p- and five 3 d-orbitals.Example 2. Determine which atom of which element has electronic formula 1 s 2 2s 2 2p 6 3s 2 3p 1 .
In order to determine which element it is, it is necessary to find out its serial number, which is equal to the total number of electrons of the atom. In this case: 2 + 2 + 6 + 2 + 1 \u003d 13. This is aluminum.After making sure that everything you need is learned, proceed to the assignments. We wish you every success.
Recommended reading:- OS Gabrielyan and others. Chemistry 11 class. M., Bustard, 2002;
- G.E. Rudzitis, F.G. Feldman. Chemistry 11 class. M., Education, 2001.
The periodic table of chemical elements is a classification of chemical elements based on certain structural features of the atoms of chemical elements. It was compiled on the basis of the Periodic Law, discovered in 1869 by D.I.Mendeleev. At that time, the Periodic Table included 63 chemical elements and was different in appearance from the modern one. Now the Periodic Table includes about one hundred and twenty chemical elements.
The periodic table is compiled in the form of a table, in which the chemical elements are arranged in a certain order: as their atomic masses grow. Now there are many types of images of the Periodic Table. The most common image is in the form of a table with the arrangement of elements from left to right.
All chemical elements in the Periodic Table are combined into periods and groups. The periodic system includes seven periods and eight groups. Periods are called horizontal rows of chemical elements in which the properties of the elements change from typical metallic to non-metallic. Vertical columns of chemical elements, which contain elements with similar chemical properties, form groups of chemical elements.
The first, second and third periods are called small because they contain a small number of elements (the first - two elements, the second and third - eight elements each). Elements of the second and third periods are called typical, their properties regularly change from a typical metal to an inert gas.
All other periods are called large (the fourth and fifth contain 18 elements each, the sixth - 32 and the seventh - 24 elements). Elements that are inside large periods at the end of each even row exhibit a special similarity in properties. These are the so-called triads: Ferum - Cobalt - Nichol, forming the family of iron, and two others: Ruthenium - Rhodium - Palladium and Osmium - Iridium - Platinum, which form the family of platinum metals (platinoids).
At the bottom of DI Mendeleev's table are the chemical elements that form the lanthanide family and the actinide family. All these elements formally belong to the third group and come after the chemical elements lanthanum (number 57) and actinium (number 89).
The periodic table of elements contains ten rows. Small periods (first, second and third) consist of one row, large periods (fourth, fifth and sixth) contain two rows each. In the seventh period there is one row.
Each big period consists of odd and even rows. Paired rows contain metal elements, odd rows change the properties of elements as in standard elements, i.e. from metallic to pronounced non-metallic.
Each group of DI Mendeleev's table consists of two subgroups: main and secondary. The main subgroups include elements of both small and large periods, that is, the main subgroups begin with either the first or second period. The side subgroups include elements of only large periods, i.e. side subgroups begin only from the fourth period.