The composition of a molecule. That is, by what atoms the molecule is formed, in what quantity, by what bonds these atoms are connected. All this determines the property of the molecule, and, accordingly, the property of the substance that these molecules form.
For example, the properties of water: transparency, fluidity, the ability to cause rust are due precisely to the presence of two hydrogen atoms and one oxygen atom.
Therefore, before proceeding to the study of the properties of molecules (that is, the properties of substances), it is necessary to consider the “building blocks” by which these molecules are formed. Understand the structure of the atom.
How is an atom arranged?
Atoms are particles that, when combined with each other, form molecules.
The atom itself is made up of positively charged nucleus (+) and negatively charged electron shell (-). In general, the atom is electrically neutral. That is, the charge of the nucleus is equal in absolute value to the charge of the electron shell.
The nucleus is formed by the following particles:
- Protons. One proton carries a +1 charge. Its mass is 1 amu (atomic mass unit). These particles are necessarily present in the nucleus.
- Neutrons. The neutron has no charge (charge = 0). Its mass is 1 amu. Neutrons may not be in the nucleus. It is not a required component of the atomic nucleus.
Thus, protons are responsible for the total charge of the nucleus. Since one neutron has a charge of +1, the charge of the nucleus is equal to the number of protons.
The electron shell, as the name implies, is formed by particles called electrons. If we compare the nucleus of an atom with a planet, then electrons are its satellites. Revolving around the nucleus (for now let's imagine that in orbits, but in fact in orbits), they form an electron shell.
- Electron is a very small particle. Its mass is so small that it is taken as 0. But the charge of an electron is -1. That is, the modulus is equal to the charge of the proton, differs in sign. Since one electron carries a charge of -1, the total charge of the electron shell is equal to the number of electrons in it.
One important consequence, since an atom is a particle that does not have a charge (the charge of the nucleus and the charge of the electron shell are equal in absolute value, but opposite in sign), that is, electrically neutral, therefore, the number of electrons in an atom is equal to the number of protons.
How do atoms of different chemical elements differ from each other?
Atoms of different chemical elements differ from each other in the charge of the nucleus (that is, the number of protons, and, consequently, the number of electrons).
How to find out the charge of the nucleus of an atom of an element? The brilliant domestic chemist D. I. Mendeleev, having discovered the periodic law, and having developed a table named after him, gave us the opportunity to do this. His discovery was far ahead of the curve. When it was not yet known about the structure of the atom, Mendeleev arranged the elements in the table in order of increasing nuclear charge.
That is, the serial number of an element in the periodic system is the charge of the nucleus of an atom of a given element. For example, oxygen has a serial number of 8, respectively, the charge of the nucleus of the oxygen atom is +8. Accordingly, the number of protons is 8, and the number of electrons is 8.
It is the electrons in the electron shell that determine the chemical properties of the atom, but more on that later.
Now let's talk about the mass.
One proton is one unit of mass, one neutron is also one unit of mass. Therefore, the sum of neutrons and protons in the nucleus is called mass number. (The electrons do not affect the mass in any way, since we neglect its mass and consider it equal to zero).
The atomic mass unit (a.m.u.) is a special physical quantity for designating small masses of particles that form atoms.
All these three atoms are atoms of one chemical element - hydrogen. Because they have the same nuclear charge.
How will they differ? These atoms have different mass numbers (due to the different number of neutrons). The first atom has a mass number of 1, the second has 2, and the third has 3.
Atoms of the same element that differ in the number of neutrons (and hence mass numbers) are called isotopes.
The presented hydrogen isotopes even have their own names:
- The first isotope (mass number 1) is called protium.
- The second isotope (mass number 2) is called deuterium.
- The third isotope (with a mass number of 3) is called tritium.
Now the next reasonable question is: why if the number of neutrons and protons in the nucleus is an integer, their mass is 1 amu, then in the periodic system the mass of an atom is a fractional number. For sulfur, for example: 32.066. 
Answer: an element has several isotopes, they differ from each other in mass numbers. Therefore, the atomic mass in the periodic table is the average value of the atomic masses of all isotopes of an element, taking into account their occurrence in nature. This mass, given in the periodic system, is called relative atomic mass.
For chemical calculations, indicators of just such an “averaged atom” are used. Atomic mass is rounded to the nearest integer.
The structure of the electron shell.
The chemical properties of an atom are determined by the structure of its electron shell. The electrons around the nucleus are not arranged anyhow. Electrons are localized in electron orbitals.
Electronic orbital- the space around the atomic nucleus, where the probability of finding an electron is greatest.
An electron has one quantum parameter called spin. If we take the classical definition from quantum mechanics, then spin is the intrinsic angular momentum of the particle. In a simplified form, this can be represented as the direction of rotation of a particle around its axis.
An electron is a particle with a half-integer spin, an electron can have either +½ or -½ spin. Conventionally, this can be represented as a clockwise and counterclockwise rotation.
No more than two electrons with opposite spins can be in one electron orbital.
The generally accepted designation of an electronic dwelling is a cell or a dash. The electron is indicated by an arrow: the up arrow is an electron with a positive spin +½, the down arrow ↓ is an electron with a negative spin -½.
An electron that is alone in an orbital is called unpaired. Two electrons in the same orbital are called paired.
Electronic orbitals are divided into four types depending on the shape: s, p, d, f. Orbitals of the same shape form a sublevel. The number of orbitals at a sublevel is determined by the number of possible locations in space.
- s orbital.
The s orbital is spherical:
In space, the s-orbital can only be located in one way:
Therefore, the s-sublevel is formed by only one s-orbital.
- p-orbital.
The p orbital is shaped like a dumbbell:
In space, the p-orbital can only be located in three ways:
Therefore, the p-sublevel is formed by three p-orbitals.
- d-orbital.
The d-orbital has a complex shape:
In space, the d-orbital can be located in five different ways. Therefore, the d-sublevel is formed by five d-orbitals.
- f-orbital
The f-orbital has an even more complex shape. In space, the f-orbital can be placed in seven different ways. Therefore, the f-sublevel is formed by seven f-orbitals.
The electron shell of an atom is like a puff pastry. It also has layers. Electrons located on different layers have different energies: on layers closer to the nucleus - less, on those far from the nucleus - more. These layers are called energy levels.
Filling of electron orbitals.
The first energy level has only the s-sublevel:
At the second energy level, there is an s-sublevel and a p-sublevel appears:
At the third energy level, there is an s-sublevel, a p-sublevel, and a d-sublevel appears:
At the fourth energy level, in principle, an f-sublevel is added. But in the school course, f-orbitals are not filled, so we can not depict the f-sublevel:
The number of energy levels in an atom of an element is period number. When filling electron orbitals, the following principles should be followed:
- Each electron tries to occupy the position in the atom where its energy will be minimal. That is, first the first energy level is filled, then the second, and so on.
To describe the structure of the electron shell, the electronic formula is also used. The electronic formula is a short one-line record of the distribution of electrons by sublevels.
- At the sublevel, each electron first fills a vacant orbital. And each has spin +½ (up arrow).
And only after there is one electron in each sublevel orbital, the next electron becomes paired - that is, it occupies an orbital that already has an electron:
- d-sublevel is filled in a special way.
The fact is that the energy of the d-sublevel is higher than the energy of the s-sublevel of the NEXT energy layer. And as we know, the electron tries to take that position in the atom, where its energy will be minimal.
Therefore, after filling the 3p sublevel, the 4s sublevel is filled first, after which the 3d sublevel is filled.
And only after the 3d sublevel is completely filled, the 4p sublevel is filled.
It is the same with the 4th energy level. After the 4p sublevel is filled, the 5s sublevel is filled next, followed by the 4d sublevel. And after it only 5p.
- And there is one more point, one rule regarding the filling of the d-sublevel.
Then there is a phenomenon called failure. In case of failure, one electron from the s-sublevel of the next energy level literally falls to the d-electron.
Ground and excited states of the atom.
The atoms whose electronic configurations we have now built are called atoms in basic condition. That is, this is a normal, natural, if you like, state.
When an atom receives energy from outside, excitation can occur.
Excitation is the transition of a paired electron to an empty orbital, within the outer energy level.
For example, for a carbon atom:
Excitation is characteristic of many atoms. This must be remembered, because excitation determines the ability of atoms to bind to each other. The main thing to remember is the condition under which excitation can occur: a paired electron and an empty orbital in the outer energy level.
There are atoms that have several excited states:
Electronic configuration of the ion.
Ions are particles that atoms and molecules turn into by gaining or losing electrons. These particles have a charge, because they either "not enough" electrons, or their excess. Positively charged ions are called cations, negative - anions.
The chlorine atom (has no charge) gains an electron. The electron has a charge of 1- (one minus), respectively, a particle is formed that has an excess negative charge. Chlorine anion:
Cl 0 + 1e → Cl –
The lithium atom (also having no charge) loses an electron. An electron has a charge of 1+ (one plus), a particle is formed, with a lack of a negative charge, that is, its charge is positive. lithium cation:
Li 0 – 1e → Li +
Turning into ions, atoms acquire such a configuration that the external energy level becomes "beautiful", that is, completely filled. This configuration is the most thermodynamically stable, so there is a reason for atoms to turn into ions.
And therefore, the atoms of the elements of the VIII-A group (the eighth group of the main subgroup), as stated in the next paragraph, are noble gases, such are chemically inactive. They have the following structure in the ground state: the outer energy level is completely filled. Other atoms, as it were, tend to acquire the configuration of these most noble gases, and therefore turn into ions and form chemical bonds.
STRUCTURE OF MULTI-ELECTRON ATOMS AND IONS
An electron in an atom exists in the form electron cloud, that is, a certain region of the nuclear space, which covers approximately 90% of the charge and mass of the electron. This region of space is called orbital. To fully characterize the state of each electron in an atom, it is necessary to indicate the values of four quantum numbers for it: chief n , orbital l , magnetic m l and spin m s .
Principal quantum number characterizes the main energy reserve of the electron and the size of the electron cloud. It can only take positive integer values between 1 and ¥. The greater the value n, the larger the size of the electron cloud. A collection of electronic states that have the same value n, is called electronic layer or energy level. The following letter designations are accepted for energy levels
At n= 1 electron energy has a minimum value E 1 = -13.6 eV. This state of the electron is called main or normal. States since n= 2, 3, 4… are called excited. The energies corresponding to them are associated with E 1 expression
When an electron moves from one energy level to another, a quantum of electromagnetic energy D is absorbed or emitted E
where with is the speed of light ( with= 3×10 8 m/s); with/ l \u003d n - radiation frequency, s -1.
Orbital(otherwise side or azimuth) the quantum number determines the moment of momentum of the electron and characterizes the shape of the electron cloud. Can take all integer values from 0 to ( n- one). Each value l corresponds to its own shape of the electron cloud: at l= 0 – spherical; l= 1 - dumbbell; l= 2 - two dumbbells intersecting at right angles.
Electrons of the same energy level having the same values l, form energy sublevels, which have the following letter designations
The energy values in the sublevels of each level are somewhat different. The number of sublevels into which the energy level is split is equal to the level number, that is, the value n.
The state of an electron corresponding to certain values n and l, is written as a combination of digital value n and letter l(for example, when n= 3 and l= 1 write 3 p).
Magnetic quantum number characterizes the spatial orientation of the electron cloud, takes all integer values from - l before +l, in total in each sublevel (2 l+ 1) values. Number of values accepted m l, indicates the number of possible positions of an electron cloud of a given type in space, that is, the number of orbitals in a sublevel. Yes, any s sublevel consists of one orbital, p- sublevel - from 3, d- sublevel - from 5, and f- sublevel - out of 7. All orbitals of the same level have the same energy and are called degenerate.
The state of an electron in an atom, characterized by the values of quantum numbers n,l and m l, is called atomic orbital(AO).
Spin quantum number characterizes the intrinsic mechanical moment of the electron associated with its rotation around its axis. It can only take two values m s= +1/2 and m s = – 1/2.
When distributing electrons in an atom over AO, several principles and rules are followed. According to minimum energy principle electrons in an atom tend to occupy first of all those AOs that correspond to the lowest value of the electron energy. The implementation of this principle is carried out on the basis of Klechkovsky rules:
with an increase in the atomic number of the element, electrons are placed on the AO sequentially as the sum ( n+l); for the same values of this sum, the orbital with a smaller value of the number is filled earlier n .
According to the Klechkovsky rule, the filling of energy levels basically corresponds to the following series: 1 s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p etc.
Degenerate orbitals of the same level are filled with electrons in accordance with Hund's (Hund's) rule:
within the energy sublevel, electrons are arranged so that their total spin is maximum.
This means that at first the electrons fill all the free orbitals of the sublevel one at a time, having identically directed spins, and only then these AOs are filled with second (paired) electrons. In accordance with Pauli principle one AO can contain no more than two electrons that differ from each other by the value m s. Thus, the maximum electronic capacity of any s-sublevel is equal to two, p- sublevel - six, d- sublevel - 10 e, a f- sublevel - 14 e.
The total number of AO at the energy level is determined by the formula
N AO = n 2 (6)
The total number of electrons in a level can be calculated from the equation
N e = 2n 2 (7)
When one or more electrons are removed from an atom, it becomes a positively charged ion. cation, whose charge is equal to the number of electrons removed. The attachment of one or more electrons to an atom leads to the formation of a negative ion - anion, whose negative charge is equal to the number of electrons received.
When a cation is formed, first of all, the atom leaves the electrons of the external energy level, since in this case the energy costs for detaching an electron will be minimal. When an anion is formed, electrons are placed on levels in accordance with the principle of minimum energy.
Valence called electrons that are located on the external energy level and individual sublevels of the second (for lanthanides and actinides - third) from the end of the electronic layer, which are not fully formed, that is, the number of electrons in the sublevel has not reached the limit value.
Elements whose atoms are filled s-orbitals belong to the family s-elements; in which is filled p sublevel, belong to the family p-elements, etc.
Example 1 The quantum numbers of the valence electrons of the E 2- ion are
Electron number n l m l m s
Determine the ordinal number of the element and name it.
Decision
The valence electronic formula of the ion E 2-: ... 3 s 2 3p one . After the removal of two extra electrons, the electronic configuration of the atom will take the form E: ... 3 s one . Add the missing electrons E:1 s 2 2s 2 2p 6 3s one . The total number of electrons (2 + 2 + 6 + 1) \u003d 11, which means this is element number 11 - sodium Na.
Example 2 Write down the complete electronic formula of the element with serial number 27. Mark its valence electrons and indicate the values of all quantum numbers for them. What electron family does this atom belong to? Write down the electronic formula of the valence sublevels of a given atom after the removal of two valence electrons.
Decision
Element with number 27 - cobalt Co. We compose its electronic formula
27 Co: 1 s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7
Valence electrons are 4 electrons s and 3 d sublevels. The quantum number values for each of the nine valence electrons are
Electron number n l m l m s
Since the sublevel is filled d, then cobalt belongs to the family d-elements.
When two electrons are detached from a cobalt atom, a Co 2+ ion is formed. The electronic formula of valence electrons Co 2+: ... 4 d 7 5s 0 .
Example 3 Write down the electronic formulas of the silicon atom in the normal and excited states.
Decision
The electronic formula of the silicon atom contains 14 electrons. In normal condition Si 14:1 s 2 2s 2 2p 6 3s 2 3p 2. When excited, one of the paired electrons 3 s-orbitals will move to sublevel 3 p and the electronic formula will take the form
Si + E® Si * : 1 s 2 2s 2 2p 6 3s 1 3p 3 .
Further excitation of the silicon atom is impossible, since all the valence electrons of the atom are unpaired.
Tasks
1. An atom of which element in the ground state has an electronic configuration of 1 s 2 2s 2 2p 6 3s 2? Determine the total number of energy levels and sublevels occupied by electrons in a given atom.
2. Using Hund's rule, distribute electrons in orbitals corresponding to the lowest energy state of atoms: manganese, nitrogen, silicon.
3. How many free f-orbitals is contained in the atoms of elements with serial numbers 59, 60, 90, 93? Using Hund's rule, distribute the electrons among the orbitals for the atoms of these elements.
4. Write the electronic formulas of the yet undiscovered elements No. 110 and No. 113 and indicate what place they will take in the periodic system.
5. An atom of an element has an electronic formula of 1 s 2 2s 2 2R 6 3s 2 3R 6. Write for it the electronic formulas of the ion E - and the conditional ion E 7+.
6. Write the electronic formulas of atoms of elements with serial numbers 21 and 23. How many free d-orbitals in the atoms of these elements? Specify the valence electrons of the elements.
7. Write down the electronic formulas of atoms and ions: Se, Ti 2+, V 3-. Label their valence electrons.
8. Write the electronic formula of an atom and name the element if the values of the quantum numbers of valence electrons are equal:
9. For elements of which periods, the electrons of the outer layer are characterized by the value of the sum ( n+l) = 5? What electronic families do these elements belong to?
10. Write down the electronic formulas of the particles: Br - , Br + , Br 5+ . Write down the quantum numbers of the valence electrons of the Br + ion.
11. Determine the number of unpaired electrons in an iridium atom. Specify the values of quantum numbers of valence electrons of this atom.
12. Write down the electronic formula of the sulfur atom, how many unpaired electrons does this atom have in the normal and excited states? What are the electronic formulas of S 2- and S 4+?
13. How many and what spatial orientations d Orbitals do you know? What is the quantum number for this?
14. Write down the full electronic formulas of atoms and ions: Zn 4-, Kr, Se 2+. Label their valence electrons.
15. Determine the serial number of the element and write down the full electronic formula of the atom, if after attaching two electrons to it, the quantum numbers of the valence sublevels are as follows:
16. Write the electronic formulas of particles: Po, Bi 3+, Mn 2-. Draw electron-graphic diagrams of their valence sublevels.
17. Write down the complete electronic formula and electronic graphic scheme of the valence sublevels of thallium and krypton atoms.
18. Determine the total number of electrons not 8 energy level.
19. How many free d-orbitals are present in the atoms of titanium and vanadium? Write down for these atoms the values of the quantum numbers of the outer layer.
20. How many values of the magnetic quantum number are possible for electrons of the energy sublevel, the orbital quantum number of which is: a) l= 3; b) l = 4?
21. Which element has three electrons in an atom, for each of which n= 3 and l= 1? What are the values of the magnetic quantum number for them? Does this atom have paired electrons?
22. Make electronic formulas of elements with serial numbers 27 and 60. Indicate the values of all quantum numbers for valence electrons of ions of these elements with charges + 1 and - 1.
23. Can configurations exist R 7 or d 12 - electrons. Why? Compose the electronic formula of an atom of an element with serial number 22 and indicate its valence electrons.
24. Write the electronic formulas of atoms of elements with serial numbers 15 and 28. What is the maximum spin R-electrons at the atoms of the first and d-electrons at the atoms of the second element.
25. An atom of which element has the following structure of the outer and penultimate electron layers 2 s2 2R 6 3s 2 3R
26. An atom of which element has the following structure of the outer and penultimate electron layers 3 s 2 3R 6 3d 3 4s 2? Write down for them the quantum numbers of valence electrons in the normal state.
27. An atom of which element has the following structure of the outer and penultimate electron layers 3 s 2 3R 6 3d 10 4s 2 4R 5 ? Write down for them the quantum numbers of valence electrons in an excited state.
28. An atom of which element has the following structure of the outer and penultimate electron layers 4 s 2 4R 6 4d 7 5s one ? Write down the full electronic formulas for them in the excited state.
29. An atom of which element has the following structure of the outer and penultimate electron layers 4 s 2 4R 6 4d 10 5s 0? Write down the full electronic formulas for them in the excited state.
30. How many free d-orbitals are present in the atoms of niobium and zirconium? Write down for these atoms the values of the quantum numbers of the outer layer.
The process of H2+ particle formation can be represented as follows:
H + H+ H2+.
Thus, one electron is located on the bonding molecular s-orbital.
The multiplicity of the bond is equal to the half-difference of the number of electrons in the bonding and loosening orbitals. Hence, the multiplicity of the bond in the H2+ particle is equal to (1 – 0):2 = 0.5. The VS method, in contrast to the MO method, does not explain the possibility of bond formation by one electron.
The hydrogen molecule has the following electronic configuration:
The H2 molecule has two bonding electrons, which means that the bond in the molecule is single.
The molecular ion H2- has an electronic configuration:
H2- [(s 1s)2(s *1s)1].
The multiplicity of the bond in H2- is (2 - 1): 2 = 0.5.
Let us now consider homonuclear molecules and ions of the second period.
The electronic configuration of the Li2 molecule is as follows:
2Li(K2s)Li2 .
The Li2 molecule contains two bonding electrons, which corresponds to a single bond.
The process of formation of the Be2 molecule can be represented as follows:
2 Be(K2s2) Be2 .
The number of bonding and loosening electrons in the Be2 molecule is the same, and since one loosening electron destroys the action of one bonding electron, the Be2 molecule in the ground state was not found.
In a nitrogen molecule, 10 valence electrons are located in orbitals. Electronic structure of the N2 molecule:
Since there are eight bonding and two loosening electrons in the N2 molecule, this molecule has a triple bond. The nitrogen molecule is diamagnetic because it does not contain unpaired electrons.
On the orbitals of the O2 molecule, 12 valence electrons are distributed, therefore, this molecule has the configuration:
Rice. 9.2. Scheme of the formation of molecular orbitals in the O2 molecule (only 2p electrons of oxygen atoms are shown)
In the O2 molecule, in accordance with Hund's rule, two electrons with parallel spins are placed one at a time in two orbitals with the same energy (Fig. 9.2). According to the VS method, the oxygen molecule does not have unpaired electrons and should have diamagnetic properties, which is inconsistent with the experimental data. The molecular orbital method confirms the paramagnetic properties of oxygen, which are due to the presence of two unpaired electrons in the oxygen molecule. The multiplicity of bonds in an oxygen molecule is (8–4):2 = 2.
Let us consider the electronic structure of the O2+ and O2- ions. In the O2+ ion, 11 electrons are placed in its orbitals, therefore, the configuration of the ion is as follows:
The multiplicity of the bond in the O2+ ion is (8–3):2 = 2.5. In the O2- ion, 13 electrons are distributed in its orbitals. This ion has the following structure:
O2-.
The multiplicity of bonds in the O2- ion is (8 - 5): 2 = 1.5. O2- and O2+ ions are paramagnetic, as they contain unpaired electrons.
The electronic configuration of the F2 molecule has the form:
The bond multiplicity in the F2 molecule is 1, since there is an excess of two bonding electrons. Since there are no unpaired electrons in the molecule, it is diamagnetic.
In the N2, O2, F2 series, the energies and bond lengths in molecules are:
An increase in the excess of binding electrons leads to an increase in the binding energy (bond strength). When passing from N2 to F2, the bond length increases, which is due to the weakening of the bond.
In the O2-, O2, O2+ series, the bond multiplicity increases, the bond energy also increases, and the bond length decreases.
1. An atom consists of a positively charged nucleus and a negatively charged electron shell. The atom is electrically neutral. The number of protons in the nucleus is equal to the number of electrons. The nucleus is made up of protons and neutrons. The relative masses of the proton and neutron are equal to 1, the proton has a charge of +1, the neutron is not charged. The nuclear charge is equal to the number of protons, the mass of the nucleus is equal to the sum of the masses of protons and neutrons. The mass of an atom consists mainly of the mass of the nucleus, since the mass of electrons is small (the mass of an electron is 1/1840 of the mass of a proton).
2. The serial number of the element is equal to the nuclear charge (the number of protons), the relative mass of the isotope of the element is equal to the number of protons and neutrons: Ar = Z + N.
3. Electrons are placed in energy levels. The number of energy levels in an atom is equal to the period number. The maximum number of electrons in an energy level is 2n 2 (n is the energy level number).
4. Electrons that are at the same energy level form different clouds (orbitals):
s - electrons form clouds of spherical shape,
p - electrons - dumbbell-shaped,
d and f - electrons have a more complex shape.
At the first energy level there is only an s-sublevel, at the second s- and p-sublevels, at the third s-, p-, d-sublevels, at the fourth s-, p-, d-, f-sublevels.
The energy sublevels have one s-orbital, three p-orbitals, five d-orbitals, seven f-orbitals. Each orbital can have one (unpaired) or two (paired) electrons. Thus, the maximum number of s-electrons at the energy level is 2, p-electrons - 6, d-electrons - 10, f-electrons - 14.
5. Energy level can be completed or incomplete. In a completed energy level, all orbitals are filled, electrons are paired.
6. The filling of energy levels follows the principle of least energy. The electron occupies the orbital with the lowest energy.
7. The electronic structure is written by an electronic formula (for example: 6 C 1s 2 2s 2 2p 2) or using quantum cells.
8. The chemical properties of an element depend on the electronic structure. The electronic structure of atoms is periodically repeated, therefore, chemical properties are periodically repeated.
9. The highest oxidation state (and highest valency) for most elements is determined by the group number.
10. The negative oxidation state of non-metals (valence in volatile hydrogen compounds of non-metals) is determined by the number of electrons missing before the completion of the external energy level, according to the formula "group number - 8".
11. Ions are formed from atoms as a result of giving or receiving electrons.
E 0 - ne \u003d E n +
E 0 + ne \u003d E n-
12. Isotopes - atoms of the same chemical element that have the same nuclear charge, but different masses. Isotopic nuclei contain the same number of protons but different numbers of neutrons.
