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PERIODIC TABLE OF ELEMENTS is a classification of chemical elements in accordance with the periodic law, which establishes a periodic change in the properties of chemical elements as their atomic mass increases, associated with an increase in the charge of the nucleus of their atoms; therefore, the charge of the nucleus of an atom coincides with the ordinal number of the element in the periodic system and is called atomic number element. The periodic system of elements is drawn up in the form of a table (periodic table of elements), in the horizontal rows of which - periods- there is a gradual change in the properties of the elements, and in the transition from one period to another - a periodic repetition of common properties; vertical columns - groups- combine elements with similar properties. The periodic system allows, without special studies, to learn about the properties of an element only on the basis of the known properties of elements neighboring in a group or period. Physical and chemical properties (aggregate state, hardness, color, valency, ionization, stability, metallicity or non-metallicity, etc.) can be predicted for an element based on the periodic table.
At the end of the 18th and beginning of the 19th centuries. chemists tried to create classifications of chemical elements in accordance with their physical and chemical properties, in particular, on the basis of the aggregate state of the element, specific gravity (density), electrical conductivity, metallicity - non-metallicity, basicity - acidity, etc.
Classifications by "atomic weight"
(i.e. by relative atomic mass).
Prout's hypothesis.
| Table 1. PERIODIC TABLE OF ELEMENTS PUBLISHED BY MENDELEEV IN 1869 (first version) |
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| Ti = 50 | Zr = 90 | ? = 180 | |||
| V=51 | Nb = 94 | Ta = 182 | |||
| cr=52 | Mo = 96 | W=186 | |||
| Mn = 55 | Rh = 104.4 | Pt = 197.4 | |||
| Fe = 56 | Ru = 104.4 | Ir = 198 | |||
| Ni = | Co = 59 | Pd = 106.6 | Os = 199 | ||
| H=1 | Cu = 63.4 | Ag = 108 | Hg = 200 | ||
| Be = 9.4 | Mg = 24 | Zn = 65.2 | CD = 112 | ||
| B=11 | Al = 27.4 | ? = 68 | Ur = 116 | Au = 197? | |
| C=12 | Si = 28 | ? = 70 | Sn = 118 | ||
| N=14 | P=31 | As = 75 | Sb = 122 | Bi = 210? | |
| O=16 | S=32 | Se = 79.4 | Te = 128? | ||
| F=19 | Cl = 35.5 | Br = 80 | I=127 | ||
| Li = 7 | Na = 23 | K = 39 | Rb = 85.4 | Cs = 133 | Tl = 204 |
| Ca=40 | Sr = 87.6 | Ba = 137 | Pb = 207 | ||
| ? = 45 | Ce = 92 | ||||
| ?Er = 56 | La = 94 | ||||
| ?Yt = 60 | Di = 95 | ||||
| ?In = 75.6 | th = 118 | ||||
| Table 2. MODIFIED MENDELEEV'S TABLE | |||||||||||||||||||||||||
| Group | I | II | III | IV | V | VI | VII | VIII | 0 | ||||||||||||||||
| Oxide or hydride formula Subgroup |
R2O | RO | R2O3 | RH4 RO 2 |
RH 3 R2O5 |
RH 2 RO 3 |
RH R2O7 |
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| Period 1 | 1 H Hydrogen 1,0079 |
2 He Helium 4,0026 |
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| Period 2 | 3 Li Lithium 6,941 |
4 Be Beryllium 9,0122 |
5 B Bor 10,81 |
6 C Carbon 12,011 |
7 N Nitrogen 14,0067 |
8 O Oxygen 15,9994 |
9 F Fluorine 18,9984 |
10 Ne Neon 20,179 |
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| Period 3 | 11 Na Sodium 22,9898 |
12 mg Magnesium 24,305 |
13 Al Aluminum 26,9815 |
14 Si Silicon 28,0855 |
15 P Phosphorus 30,9738 |
16 S Sulfur 32,06 |
17 Cl Chlorine 35,453 |
18 Ar Argon 39,948 |
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| Period 4 | 19 K Potassium 39,0983 29 Cu Copper 63,546 |
20 Ca Calcium 40,08 30 Zn Zinc 65,39 |
21 sc Scandium 44,9559 31 Ga Gallium 69,72 |
22 Ti Titanium 47,88 32 Ge Germanium 72,59 |
23 V Vanadium 50,9415 33 As Arsenic 74,9216 |
24 Cr Chromium 51,996 34 Se Selenium 78,96 |
25 Mn Manganese 54,9380 35 Br Bromine 79,904 |
26 Fe Iron 55,847 |
27 co Cobalt 58,9332 |
28 Ni Nickel 58,69 |
36 |
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| Period 5 | 37 Rb Rubidium 85,4678 47 Ag Silver 107,868 |
38 Sr Strontium 87,62 48 CD Cadmium 112,41 |
39 Y Yttrium 88,9059 49 In Indium 114,82 |
40 Zr Zirconium 91,22 50 sn Tin 118,69 |
41 Nb Niobium 92,9064 51 Sb Antimony 121,75 |
42 Mo Molybdenum 95,94 52 Te Tellurium 127,60 |
43 Tc Technetium 53 I iodine 126,9044 |
44 Ru Ruthenium 101,07 |
45 Rh Rhodium 102,9055 |
46 Pd Palladium 106,4 |
54 |
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| Period 6 | 55 Cs Cesium 132,9054 79 Au Gold 196,9665 |
56 Ba Barium 137,33 80 hg Mercury 200,59 |
57* La Lanthanum 138,9055 81 Tl Thallium 204,38 |
72 hf Hafnium 178,49 82 Pb Lead 207,21 |
73 Ta Tantalum 180,9479 83 Bi Bismuth 208,9804 |
74 W Tungsten 183,85 84 Po Polonium |
75 Re Rhenium 186,207 85 At Astatine |
76 Os Osmium 190,2 |
77 Ir Iridium 192,2 |
78 Pt Platinum 195,08 |
86 |
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| Period 7 | 87 Fr France |
88 Ra Radium 226,0254 |
89** AC Actinium 227,028 |
104 | 105 | 106 | 107 | 108 | 109 | ||||||||||||||||
| * | 58 Ce 140,12 |
59 Pr 140,9077 |
60 Nd 144,24 |
61 Pm |
62 sm 150,36 |
63 Eu 151,96 |
64 Gd 157,25 |
65 Tb 158,9254 |
66 Dy 162,50 |
67 Ho 164,9304 |
68 Er 167,26 |
69 Tm 168,9342 |
70 Yb 173,04 |
71 Lu 174,967 |
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| ** | 90 Th 232,0381 |
91 Pa 231,0359 |
92 U 238,0289 |
93 Np 237,0482 |
94 Pu |
95 Am |
96 cm |
97 bk |
98 cf |
99 Es |
100 fm |
101 md |
102 no |
103 lr |
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| * Lanthanides: cerium, praseodymium, neodymium, promethium, samarium, europium, gadolinium, terbium, dysprosium, holmium, erbium, thulium, ytterbium, lutetium. ** Actinides: thorium, protactinium, uranium, neptunium, plutonium, americium, curium, berkelium, californium, einsteinium, fermium, mendelevium, nobelium, lawrencium. Note. The atomic number is indicated above the element symbol, the atomic mass is indicated below the element symbol. The value in brackets is the mass number of the longest-lived isotope. |
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Periods.
In this table, Mendeleev arranged the elements in horizontal rows - periods. The table starts with a very short period containing only hydrogen and helium. The next two short periods contain 8 elements each. Then there are four long periods. All periods except the first start with an alkali metal (Li, Na, K, Rb, Cs) and all periods end with a noble gas. In the 6th period there is a series of 14 elements - lanthanides, which formally has no place in the table and is usually placed under the table. Another similar series - actinides - is in the 7th period. This series includes elements produced in the laboratory, such as by bombarding uranium with subatomic particles, and is also placed below the lanthanides below the table.
Groups and subgroups.
When the periods are arranged one below the other, the elements are arranged in columns, forming groups numbered 0, I, II, ..., VIII. The elements within each group are expected to exhibit similar general chemical properties. Even greater similarity is observed for elements in subgroups (A and B), which are formed from elements of all groups except 0 and VIII. Subgroup A is called the main subgroup, and B is called the secondary subgroup. Some families have names, such as alkali metals (Group IA), alkaline earth metals (Group IIA), halogens (Group VIIA), and noble gases (Group 0). Group VIII contains the transition metals Fe, Co, and Ni; Ru, Rh and Pd; Os, Ir and Pt. Being in the middle of long periods, these elements are more similar to each other than to the elements before and after them. In several cases, the order of increase in atomic weights (more precisely, atomic masses) is violated, for example, in pairs of tellurium and iodine, argon and potassium. This "violation" is necessary to maintain the similarity of elements in subgroups.
Metals, non-metals.
The diagonal from hydrogen to radon roughly divides all elements into metals and non-metals, while non-metals are above the diagonal. (Non-metals include 22 elements - H, B, C, Si, N, P, As, O, S, Se, Te, halogens and inert gases, metals - all other elements.) Along this line are elements that have some properties of metals and non-metals (metalloids are an obsolete name for such elements). When considering properties by subgroups from top to bottom, an increase in metallic properties and a weakening of non-metallic properties are observed.
Valence.
The most general definition of the valency of an element is the ability of its atoms to combine with other atoms in certain ratios. Sometimes the valency of an element is replaced by the concept of the oxidation state (s.o.) close to it. The oxidation state corresponds to the charge that an atom would acquire if all the electron pairs of its chemical bonds were shifted towards more electronegative atoms. In any period, from left to right, there is an increase in the positive oxidation state of the elements. Elements of group I have s.d. equal to +1 and the oxide formula R 2 O, elements of group II - respectively +2 and RO, etc. Elements with negative s.d. are in groups V, VI and VII; it is believed that carbon and silicon, which are in group IV, do not have a negative oxidation state. Halogens having an oxidation state of –1 form compounds with hydrogen of composition RH. In general, the positive oxidation state of the elements corresponds to the group number, and the negative one is equal to the difference of eight minus the group number. From the table it is impossible to determine the presence or absence of other oxidation states.
The physical meaning of the atomic number.
A true understanding of the periodic table is possible only on the basis of modern ideas about the structure of the atom. The atomic number of an element in the periodic table is much more important than its atomic weight (i.e., relative atomic mass) for understanding chemical properties.
The structure of the atom.
In 1913, N. Bohr used the nuclear model of the structure of the atom to explain the spectrum of the hydrogen atom, the lightest and therefore the simplest atom. Bohr suggested that the hydrogen atom consists of one proton, which makes up the nucleus of the atom, and one electron, which revolves around the nucleus.
Definition of the concept of atomic number.
In 1913, A. van den Broek suggested that the atomic number of an element - its atomic number - should be identified with the number of electrons revolving around the nucleus of a neutral atom, and with the positive charge of the atomic nucleus in units of electron charge. However, it was necessary to experimentally confirm the identity of the charge of the atom and the atomic number. Bohr further postulated that the characteristic X-ray emission of an element should follow the same law as the spectrum of hydrogen. Thus, if the atomic number Z is identified with the charge of the nucleus in units of electron charge, then the frequencies (wavelengths) of the corresponding lines in the X-ray spectra of various elements should be proportional to Z 2 , the square of the element's atomic number.
In 1913-1914, G. Moseley, studying the characteristic X-ray radiation of atoms of various elements, received a brilliant confirmation of Bohr's hypothesis. Moseley's work thus confirmed van den Broek's assumption that the atomic number of an element is identical with the charge of its nucleus; atomic number, not atomic mass, is the true basis for determining the chemical properties of an element.
Periodicity and atomic structure.
Bohr's quantum theory of the structure of the atom developed over the two decades after 1913. Bohr's proposed "quantum number" became one of the four quantum numbers needed to characterize the energy state of an electron. In 1925, W. Pauli formulated his famous "principle of prohibition" (Pauli principle), according to which there cannot be two electrons in an atom, in which all quantum numbers would be the same. When this principle was applied to the electronic configurations of atoms, the periodic table acquired a physical basis. Since the atomic number Z, i.e. If the positive charge of the nucleus of an atom increases, then the number of electrons must also increase in order to maintain the electroneutrality of the atom. These electrons determine the chemical "behavior" of the atom. According to the Pauli principle, as the value of the quantum number increases, the electrons fill the electron layers (shells) starting from those closest to the nucleus. The completed layer, which is filled with all electrons according to the Pauli principle, is the most stable. Therefore, noble gases such as helium and argon, which have fully completed electronic structures, are resistant to any chemical attack.
Electronic configurations.
The following table lists the possible numbers of electrons for various energy states. Principal quantum number n= 1, 2, 3,... characterizes the energy level of electrons (the 1st level is located closer to the nucleus). Orbital quantum number l = 0, 1, 2,..., n– 1 characterizes the orbital angular momentum. The orbital quantum number is always less than the main quantum number, and its maximum value is equal to the main quantum number minus 1. Each value l corresponds to a certain type of orbital - s, p, d, f... (this designation comes from the spectroscopic nomenclature of the 18th century, when various series of observed spectral lines were called s harp, p rincipal, d diffuse and f undamental).
| Table 3. NUMBER OF ELECTRONS IN VARIOUS ENERGY STATES OF THE ATOM | |||
| Principal quantum number | Orbital quantum number | The number of electrons on the shell | Energy state designation (orbital type) |
| 1 | 0 | 2 | 1s |
| 2 | 0 | 2 | 2s |
| 1 | 6 | 2p | |
| 3 | 0 | 2 | 3s |
| 1 | 6 | 3p | |
| 2 | 10 | 3d | |
| 4 | 0 | 2 | 4s |
| 1 | 6 | 4p | |
| 2 | 10 | 4d | |
| 3 | 14 | 4f | |
| 5 | 0 | 2 | 5s |
| 1 | 6 | 5p | |
| 2 | 10 | 5d | |
| 5 | 14 | 5f | |
| 4 | 18 | 5g | |
| 6 | 0 | 2 | 6s |
| 1 | 6 | 6p | |
| 2 | 10 | 6d | |
| ... | ... | ... | ... |
| 7 | 0 | 2 | 7s |
Short and long periods.
The lowest fully completed electron shell (orbital) is denoted 1 s and is realized in helium. Next levels - 2 s and 2 p- correspond to the building-up of the shells of atoms of the elements of the 2nd period and, with full building-up, for neon, contain a total of 8 electrons. As the values of the principal quantum number increase, the energy state of the lowest orbital number for the larger principal may be lower than the energy state of the highest orbital quantum number corresponding to the smaller principal. So, energy state 3 d higher than 4 s, so the elements of the 3rd period are built up 3 s- and 3 p-orbitals, ending with the formation of a stable structure of the noble gas argon. Next comes the sequential building 4 s-, 3d- and 4 p-orbitals for elements of the 4th period, up to the completion of the outer stable electron shell of 18 electrons for krypton. This leads to the appearance of the first long period. Similarly, building 5 s-, 4d- and 5 p-orbitals of atoms of the elements of the 5th (i.e. the second long) period, ending with the electronic structure of xenon.
Lanthanides and actinides.
Sequential filling with electrons 6 s-, 4f-, 5d- and 6 p-orbitals of the elements of the 6th (i.e. the third long) period leads to the appearance of new 32 electrons, which form the structure of the last element of this period - radon. Starting with the 57th element, lanthanum, 14 elements are sequentially arranged, differing little in chemical properties. They form a series of lanthanides, or rare earth elements, in which 4 f-shell containing 14 electrons.
The series of actinides, which is located behind the actinium (atomic number 89), is characterized by building up 5 f- shells; it also includes 14 elements that are very similar in chemical properties. The element with atomic number 104 (rutherfordium), which follows the last of the actinides, already differs in chemical properties: it is an analogue of hafnium. The following names are accepted for the elements after rutherfordium: 105 - dubnium (Db), 106 - seaborgium (Sg), 107 - bohrium (Bh), 108 - hassium (Hs), 109 - meitnerium (Mt).
Application of the periodic table.
Knowledge of the periodic table allows the chemist to predict with a certain degree of accuracy the properties of any element before he starts working with it. Metallurgists, for example, consider the periodic table useful for creating new alloys, since, using the periodic table, one of the metals of the alloy can be replaced by choosing a replacement for it among its neighbors in the table so that, with a certain degree of probability, there will be no significant change in the properties formed from them. alloy.
1. Specify the name of the element, its designation. Determine the element's serial number, period number, group, subgroup. Indicate the physical meaning of the system parameters - serial number, period number, group number. Justify the position in the subgroup.
2. Indicate the number of electrons, protons and neutrons in an atom of an element, nuclear charge, mass number.
3. Make a complete electronic formula of the element, determine the electronic family, assign a simple substance to the class of metals or non-metals.
4. Draw graphically the electronic structure of the element (or the last two levels).
5. Graphically depict all possible valence states.
6. Specify the number and type of valence electrons.
7. List all possible valencies and oxidation states.
8. Write the formulas of oxides and hydroxides for all valence states. Indicate their chemical nature (confirm the answer with the equations of the corresponding reactions).
9. Give the formula of a hydrogen compound.
10. Name the scope of this element
Solution. Scandium corresponds to the element with the atomic number 21 in the PSE.
1. The element is in the IV period. The period number means the number of energy levels in the atom of this element, it has 4 of them. Scandium is located in the 3rd group - on the outer level of the 3rd electron; in the side group. Therefore, its valence electrons are in the 4s and 3d sublevels. The serial number numerically coincides with the charge of the nucleus of an atom.
2. The charge of the nucleus of the scandium atom is +21.
The number of protons and electrons is 21 each.
The number of neutrons A–Z = 45 – 21 = 24.
The total composition of the atom: ( 
).
3. Full electronic formula of scandium:
1s 2 2s 2 2p 6 3s 2 3p 6 3d 1 4s 2 .
Electron family: d-element, as in the process of filling
d-orbitals. The electronic structure of the atom ends with s-electrons, so scandium exhibits metallic properties; simple substance - metal.
4. Electronic graphic configuration looks like:
5. Possible valence states due to the number of unpaired electrons:
- in basic condition:
– in scandium in an excited state, an electron from the 4s orbital will pass to a free 4p orbital, one unpaired d-electron increases the valence capabilities of scandium.
Sc has three valence electrons in the excited state.
6. Possible valencies in this case are determined by the number of unpaired electrons: 1, 2, 3 (or I, II, III). Possible oxidation states (reflecting the number of displaced electrons) +1, +2, +3 (since scandium is a metal).
7. The most characteristic and stable valency III, oxidation state +3. The presence of only one electron in the d state is responsible for the low stability of the 3d 1 4s 2 configuration.
Scandium and its analogs, unlike other d-elements, exhibit a constant oxidation state of +3, this is the highest oxidation state and corresponds to the group number.
8. Formulas of oxides and their chemical nature:
form of higher oxide - (amphoteric);
hydroxide formulas: – amphoteric.
Reaction equations confirming the amphoteric nature of oxides and hydroxides:
(scandate of lithium),
(scandium chloride),
( potassium hexahydroxoscandiate (III) ),
(scandium sulfate).
9. It does not form compounds with hydrogen, since it is in the side subgroup and is a d-element.
10. Scandium compounds are used in semiconductor technology.
Example 2 Which of the two elements, manganese or bromine, has more pronounced metallic properties?
Solution. These elements are in the fourth period. We write down their electronic formulas:
Manganese is a d-element, i.e. an element of a side subgroup, and bromine is
p-element of the main subgroup of the same group. At the outer electronic level, the manganese atom has only two electrons, while the bromine atom has seven. The radius of the manganese atom is less than the radius of the bromine atom with the same number of electron shells.
A common pattern for all groups containing p- and d-elements is the predominance of metallic properties in d-elements.
Thus, the metallic properties of manganese are more pronounced than those of bromine.
Having studied the properties of elements arranged in a row in ascending order of their atomic masses, the great Russian scientist D.I. Mendeleev in 1869 derived the law of periodicity:
the properties of the elements, and therefore the properties of the simple and complex bodies formed by them, are in a periodic dependence on the magnitude of the atomic weights of the elements.
modern formulation of Mendeleev's periodic law:
The properties of chemical elements, as well as the forms and properties of compounds of elements, are in a periodic dependence on the charge of their nuclei.
The number of protons in the nucleus determines the value of the positive charge of the nucleus and, accordingly, the serial number Z of the element in the periodic system. The total number of protons and neutrons is called mass number A, it is approximately equal to the mass of the nucleus. So the number of neutrons (N) in the kernel can be found by the formula:
N = A - Z.
Electronic configuration- the formula for the arrangement of electrons in various electron shells of an atom-chemical element
Or molecules.
17. Quantum numbers and order of filling energy levels and orbitals in atoms. Rules of Klechkovsky
The order of distribution of electrons over energy levels and sublevels in the shell of an atom is called its electronic configuration. The state of each electron in an atom is determined by four quantum numbers:
1. Principal quantum number n characterizes to the greatest extent the energy of an electron in an atom. n = 1, 2, 3….. The electron has the lowest energy at n = 1, while it is closest to the atomic nucleus.
2. Orbital (side, azimuthal) quantum number l determines the shape of the electron cloud and, to a small extent, its energy. For each value of the principal quantum number n, the orbital quantum number can take zero and a number of integer values: l = 0…(n-1)
The states of an electron characterized by different values of l are usually called the energy sublevels of an electron in an atom. Each sublevel is designated by a certain letter, it corresponds to a certain form of the electron cloud (orbital).
3. Magnetic quantum number m l determines the possible orientations of the electron cloud in space. The number of such orientations is determined by the number of values that the magnetic quantum number can take:
m l = -l, …0,…+l
The number of such values for a specific l: 2l+1
Respectively: for s-electrons: 2·0 +1=1 (a spherical orbital can be oriented in only one way);
4. Spin quantum number m s o reflects the presence of an intrinsic momentum of the electron.
The spin quantum number can only have two values: m s = +1/2 or –1/2
Distribution of electrons in multielectron atoms takes place according to three principles:
Pauli principle
An atom cannot have electrons that have the same set of all four quantum numbers.
2. Hund's rule(tram rule)
In the most stable state of the atom, electrons are located within the electronic sublevel so that their total spin is maximum. Similar to the procedure for filling double seats in an empty tram approaching the stop - first, people who do not know each other sit down on double seats (and electrons in orbitals) one by one, and only when the empty double seats run out in two.
The principle of minimum energy (Rules of V.M. Klechkovsky, 1954)
1) With an increase in the charge of the nucleus of an atom, the successive filling of electron orbitals occurs from orbitals with a smaller value of the sum of the principal and orbital fifth numbers (n + l) to orbitals with a larger value of this sum.
2) For the same values of the sum (n + l), the filling of the orbitals occurs sequentially in the direction of increasing the value of the principal quantum number.
18. Methods for modeling chemical bonds: the method of valence bonds and the method of molecular orbitals.
Valence bond method
The simplest is the method of valence bonds (BC), proposed in 1916 by the American physical chemist Lewis.
The method of valence bonds considers a chemical bond as a result of the attraction of the nuclei of two atoms to one or more electron pairs common to them. Such a two-electron and two-center bond, localized between two atoms, is called covalent.
In principle, two mechanisms for the formation of a covalent bond are possible:
1. Pairing of electrons of two atoms under the condition of opposite orientation of their spins;
2. Donor-acceptor interaction, in which a ready electron pair of one of the atoms (donor) becomes common in the presence of an energetically favorable free orbital of another atom (acceptor).
IV - VII - big periods, because consist of two rows (even and odd) of elements.
In even rows of large periods are typical metals. The odd series begins with a metal, then the metallic properties weaken and non-metallic properties increase, the period ends with an inert gas.
Group is a vertical row of chem. elements combined by chem. properties.
Group
main subgroup secondary subgroup
The main subgroup includes The secondary subgroup includes
elements of both small and large elements of only large periods.
periods.
H, Li, Na, K, Rb, Cs, Fr Cu, Ag, Au
small big big
For elements combined in the same group, the following patterns are characteristic:
1. Highest valency of elements in compounds with oxygen(with a few exceptions) corresponds to the group number.
Elements of secondary subgroups may also exhibit another higher valency. For example, Cu - an element of group I of the side subgroup - forms oxide Cu 2 O. However, the most common are compounds of divalent copper.
2. In the main subgroups(top down) with an increase in atomic masses, the metallic properties of the elements increase and the non-metallic ones weaken.
The structure of the atom.
For a long time, science was dominated by the opinion that atoms are indivisible, i.e. do not contain simpler components.
However, at the end of the 19th century, a number of facts were established that testified to the complex composition of atoms and the possibility of their mutual transformations.
Atoms are complex formations built from smaller structural units.
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ē - electron - outside the nucleus
For chemistry, the structure of the electron shell of the atom is of great interest. Under electron shell understand the totality of all electrons in an atom. The number of electrons in an atom is equal to the number of protons, i.e. the atomic number of the element, since the atom is electrically neutral.
The most important characteristic of an electron is the energy of its bond with an atom. Electrons with similar energy values form a single electronic layer.
Each chem. element in the periodic table was numbered.
The number that each element receives is called serial number.
The physical meaning of the serial number:
1. What is the serial number of the element, such is the charge of the nucleus of the atom.
2. The same number of electrons revolve around the nucleus.
Z = p + Z - element number
n 0 \u003d A - Z
n 0 \u003d A - p + A - atomic mass of the element
n 0 \u003d A - ē
For example Li.
The physical meaning of the period number.
In what period is the element, how many electron shells (layers) it will have.
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Not +2 | |
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Determination of the maximum number of electrons in one electron shell.
Periodic law of D.I Mendeleev.
The properties of chemical elements, and therefore the properties of the simple and complex bodies they form, are in a periodic dependence on the magnitude of the atomic weight.
The physical meaning of the periodic law.
The physical meaning of the periodic law lies in the periodic change in the properties of elements, as a result of periodically repeating e-th shells of atoms, with a successive increase in n.
The modern formulation of D.I. Mendeleev's PZ.
The property of chemical elements, as well as the property of the simple or complex substances formed by them, is in a periodic dependence on the magnitude of the charge of the nuclei of their atoms.
Periodic system of elements.
Periodic system - a system of classifications of chemical elements, created on the basis of the periodic law. Periodic system - establishes relationships between chemical elements reflecting their similarities and differences.
Periodic table (there are two types: short and long) of elements.
The Periodic Table of the Elements is a graphical representation of the Periodic Table of the Elements, consists of 7 periods and 8 groups.
Question 10
Periodic system and structure of electron shells of atoms of elements.
Later it was found that not only the serial number of the element has a deep physical meaning, but also other concepts previously considered earlier also gradually acquired a physical meaning. For example, the group number, indicating the highest valency of the element, thereby reveals the maximum number of electrons of an atom of a particular element that can participate in the formation of a chemical bond.
The period number, in turn, turned out to be related to the number of energy levels present in the electron shell of an atom of an element of a given period.
Thus, for example, the "coordinates" of tin Sn (serial number 50, period 5, main subgroup of group IV) mean that there are 50 electrons in the tin atom, they are distributed over 5 energy levels, only 4 electrons are valence.
The physical meaning of finding elements in subgroups of various categories is extremely important. It turns out that for elements located in subgroups of category I, the next (last) electron is located on s-sublevel external level. These elements belong to the electronic family. For atoms of elements located in subgroups of category II, the next electron is located on p-sublevel external level. These are the elements of the “p” electronic family. Thus, the next 50th electron of tin atoms is located on the p-sublevel of the outer, i.e., 5th energy level.
For atoms of elements of subgroups of category III, the next electron is located on d-sublevel, but already before the external level, these are elements of the electronic family "d". For lanthanide and actinide atoms, the next electron is located on the f-sublevel, before the external level. These are the elements of the electronic family "f".
It is no coincidence, therefore, that the numbers of subgroups of these 4 categories noted above, that is, 2-6-10-14, coincide with the maximum numbers of electrons in the s-p-d-f sublevels.
But it turns out that it is possible to solve the problem of the order of filling the electron shell and derive an electronic formula for an atom of any element and on the basis of the periodic system, which clearly indicates the level and sublevel of each successive electron. The periodic system also indicates the placement of elements one after another into periods, groups, subgroups and the distribution of their electrons by levels and sublevels, because each element has its own, characterizing its last electron. As an example, let us analyze the compilation of an electronic formula for the atom of the element zirconium (Zr). The periodic system gives the indicators and "coordinates" of this element: serial number 40, period 5, group IV, side subgroup. First conclusions: a) all 40 electrons, b) these 40 electrons are distributed over five energy levels; c) out of 40 electrons only 4 are valence, d) the next 40th electron entered the d-sublevel before the outer, i.e. the fourth energy level.Similar conclusions can be drawn about each of the 39 elements preceding zirconium, only the indicators and coordinates will be different each time.
