It is written in the form of so-called electronic formulas. In electronic formulas, the letters s, p, d, f denote the energy sublevels of electrons; numbers in front of letters mean energy level, in which a given electron is located, and the index at the top right is the number of electrons in a given sublevel. To compose the electronic formula of an atom of any element, it is enough to know the number of this element in the periodic table and follow the basic principles that govern the distribution of electrons in the atom.
The structure of the electron shell of an atom can also be depicted in the form of a diagram of the arrangement of electrons in energy cells.
For iron atoms, this scheme has the following form:
This diagram clearly shows the implementation of Hund's rule. At the 3d sublevel maximum amount, cells (four) are filled with unpaired electrons. The image of the structure of the electron shell in an atom in the form of electronic formulas and in the form of diagrams does not clearly reflect wave properties electron.
The wording of the periodic law as amended YES. Mendeleev : the properties of simple bodies, as well as the forms and properties of compounds of elements, are in a periodic dependence on the magnitude of the atomic weights of the elements.
Modern formulation of the Periodic Law: the properties of elements, as well as the forms and properties of their compounds, are periodically dependent on the magnitude of the charge of the nucleus of their atoms.
Thus, the positive charge of the nucleus (not atomic mass) turned out to be a more accurate argument on which the properties of elements and their compounds depend
Valence- This is the number of chemical bonds by which one atom is connected to another.
The valence capabilities of an atom are determined by the number of unpaired electrons and the presence of free atomic orbitals at the outer level. The structure of the outer energy levels of atoms of chemical elements mainly determines the properties of their atoms. Therefore, these levels are called valence levels. Electrons of these levels, and sometimes of pre-external levels, can take part in the formation of chemical bonds. Such electrons are also called valence electrons.
Stoichiometric valencechemical element- this is the number of equivalents that a given atom can attach to itself, or the number of equivalents in an atom.
Equivalents are determined by the number of attached or substituted hydrogen atoms, so the stoichiometric valence is equal to the number of hydrogen atoms with which a given atom interacts. But not all elements interact freely, but almost all of them interact with oxygen, so stoichiometric valence can be defined as twice the number of attached oxygen atoms.
For example, the stoichiometric valence of sulfur in hydrogen sulfide H 2 S is 2, in oxide SO 2 - 4, in oxide SO 3 -6.
When determining the stoichiometric valence of an element using the formula of a binary compound, one should be guided by the rule: the total valence of all atoms of one element must be equal to the total valence of all atoms of another element.
Oxidation state Also characterizes the composition of the substance and is equal to the stoichiometric valency with a plus sign (for a metal or a more electropositive element in the molecule) or minus.
1. In simple substances, the oxidation state of elements is zero.
2. The oxidation state of fluorine in all compounds is -1. The remaining halogens (chlorine, bromine, iodine) with metals, hydrogen and other more electropositive elements also have an oxidation state of -1, but in compounds with more electronegative elements they have positive values oxidation states.
3. Oxygen in compounds has an oxidation state of -2; the exceptions are hydrogen peroxide H 2 O 2 and its derivatives (Na 2 O 2, BaO 2, etc., in which oxygen has an oxidation state of -1, as well as oxygen fluoride OF 2, in which the oxidation state of oxygen is +2.
4. Alkaline elements (Li, Na, K, etc.) and elements main subgroup the second group of the Periodic Table (Be, Mg, Ca, etc.) always have an oxidation state equal to the group number, that is, +1 and +2, respectively.
5. All elements of the third group, except thallium, have a constant oxidation state equal to the group number, i.e. +3.
6. The highest oxidation state of an element is equal to the group number of the Periodic Table, and the lowest is the difference: group number is 8. For example, the highest oxidation state of nitrogen (it is located in the fifth group) is +5 (in nitric acid and its salts), and the lowest is equal to -3 (in ammonia and ammonium salts).
7. The oxidation states of the elements in a compound cancel each other out so that their sum for all atoms in a molecule or neutral formula unit is zero, and for an ion it is its charge.
These rules can be used to determine the unknown oxidation state of an element in a compound if the oxidation states of the others are known, and to construct formulas for multielement compounds.
Oxidation state (oxidation number) — an auxiliary conventional value for recording the processes of oxidation, reduction and redox reactions.
Concept oxidation state often used in inorganic chemistry instead of the concept valence. The oxidation state of an atom is equal to the numerical value electric charge, assigned to an atom under the assumption that the bonding electron pairs are entirely biased towards more electronegative atoms (that is, under the assumption that the compound consists only of ions).
The oxidation number corresponds to the number of electrons that must be added to a positive ion to reduce it to a neutral atom, or subtracted from a negative ion to oxidize it to a neutral atom:
Al 3+ + 3e − → Al
S 2− → S + 2e − (S 2− − 2e − → S)
The properties of elements, depending on the structure of the electron shell of the atom, change according to periods and groups periodic table. Since in a series of analogue elements the electronic structures are only similar, but not identical, then when moving from one element in the group to another, not a simple repetition of properties is observed for them, but their more or less clearly expressed natural change.
The chemical nature of an element is determined by the ability of its atom to lose or gain electrons. This ability is quantified by the values of ionization energies and electron affinities.
Ionization energy (E and) is the minimum amount of energy required for the abstraction and complete removal of an electron from an atom in the gas phase at T = 0
K without transfer to the liberated electron kinetic energy with the transformation of the atom into a positively charged ion: E + Ei = E+ + e-. Ionization energy is a positive quantity and has smallest values for alkali metal atoms and the largest for noble (inert) gas atoms.
Electron affinity (Ee) is the energy released or absorbed when an electron is added to an atom in the gas phase at T = 0
K with the transformation of an atom into a negatively charged ion without transferring kinetic energy to the particle:
E + e- = E- + Ee.
The halogens, especially fluorine, have the maximum electron affinity (Ee = -328 kJ/mol).
The values of Ei and Ee are expressed in kilojoules per mole (kJ/mol) or in electron volts per atom (eV).
The ability of a bonded atom to shift electrons of chemical bonds towards itself, increasing the electron density around itself is called electronegativity.
This concept was introduced into science by L. Pauling. Electronegativitydenoted by the symbol ÷ and characterizes the tendency of a given atom to add electrons when it forms a chemical bond.
According to R. Maliken, the electronegativity of an atom is estimated by half the sum of the ionization energies and electron affinities of free atoms = (Ee + Ei)/2
In the periods there is a general tendency for the ionization energy and electronegativity to increase with increasing charge of the atomic nucleus; in groups these values increase serial number elements decrease.
It should be emphasized that an element cannot be assigned a constant electronegativity value, since it depends on many factors, in particular on the valence state of the element, the type of compound in which it is included, and the number and type of neighboring atoms.
Atomic and ionic radii. The sizes of atoms and ions are determined by the sizes of the electron shell. According to quantum mechanical concepts, the electron shell does not have strictly defined boundaries. Therefore, the radius of a free atom or ion can be taken as theoretically calculated distance from the nucleus to the position of the main maximum of the density of the outer electron clouds. This distance is called the orbital radius. In practice, the radii of atoms and ions in compounds are usually used, calculated based on experimental data. In this case, covalent and metallic radii of atoms are distinguished.
The dependence of atomic and ionic radii on the charge of the nucleus of an element’s atom is periodic in nature. In periods, as the atomic number increases, the radii tend to decrease. The greatest decrease is typical for elements of short periods, since their outer electronic level is filled. In large periods in the families of d- and f-elements, this change is less sharp, since in them the filling of electrons occurs in the pre-external layer. In subgroups, the radii of atoms and ions of the same type generally increase.
The periodic table of elements is clear example manifestations of various kinds of periodicity in the properties of elements, which is observed horizontally (in a period from left to right), vertically (in a group, for example, from top to bottom), diagonally, i.e. some property of the atom increases or decreases, but the periodicity remains.
In the period from left to right (→), the oxidizing and non-metallic properties of the elements increase, and the reducing and metallic properties decrease. So, of all the elements of period 3, sodium will be the most active metal and the strongest reducing agent, and chlorine will be the strongest oxidizing agent.
Chemical bond- This is the mutual connection of atoms in a molecule, or crystal lattice, as a result of the action of electrical forces of attraction between the atoms.
This is the interaction of all electrons and all nuclei, leading to the formation of a stable, polyatomic system (radical, molecular ion, molecule, crystal).
Chemical bonds are carried out by valence electrons. According to modern concepts, a chemical bond is of an electronic nature, but it is carried out in different ways. Therefore, there are three main types of chemical bonds: covalent, ionic, metallic.Arises between molecules hydrogen bond, and happen van der Waals interactions.
The main characteristics of a chemical bond include:
- connection length - This is the internuclear distance between chemically bonded atoms.
It depends on the nature of the interacting atoms and the multiplicity of the bond. As the multiplicity increases, the bond length decreases and, consequently, its strength increases;
- the multiplicity of the bond is determined by the number of electron pairs connecting two atoms. As the multiplicity increases, the binding energy increases;
- connection angle- the angle between imaginary straight lines passing through the nuclei of two chemically interconnected neighboring atoms;
Bond energy E SV - this is the energy that is released during the formation of a given bond and spent on its breaking, kJ/mol.
Covalent bond - A chemical bond formed by sharing a pair of electrons between two atoms.
The explanation of the chemical bond by the emergence of shared electron pairs between atoms formed the basis of the spin theory of valency, the tool of which is valence bond method (MVS) , discovered by Lewis in 1916. For a quantum mechanical description of chemical bonds and the structure of molecules, another method is used - molecular orbital method (MMO) .
Valence bond method
Basic principles of chemical bond formation using MBC:
1. A chemical bond is formed by valence (unpaired) electrons.
2. Electrons with antiparallel spins belonging to two different atoms become common.
3. A chemical bond is formed only if, when two or more atoms approach each other, the total energy of the system decreases.
4. The main forces acting in a molecule are of electrical, Coulomb origin.
5. The stronger the connection, the more the interacting electron clouds overlap.
There are two mechanisms for the formation of covalent bonds:
Exchange mechanism. A bond is formed by sharing the valence electrons of two neutral atoms. Each atom contributes one unpaired electron to a common electron pair:
Rice. 7. Exchange mechanism for the formation of covalent bonds: A- non-polar; b- polar
Donor-acceptor mechanism. One atom (donor) provides an electron pair, and the other atom (acceptor) provides an empty orbital for that pair.
connections, educated according to the donor-acceptor mechanism, belong to complex compounds
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Rice. 8. Donor-acceptor mechanism of covalent bond formation
A covalent bond has certain characteristics.
Saturability - property of atoms to form strictly certain number covalent bonds. Due to the saturation of bonds, molecules have a certain composition.
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Directivity - t . e. the bond is formed in the direction maximum overlap electron clouds . With respect to the line connecting the centers of the atoms forming the bond, they distinguish: σ and π (Fig. 9): σ-bond - formed by overlapping AO along the line connecting the centers of interacting atoms; A π bond is a bond that occurs in the direction of an axis perpendicular to the straight line connecting the nuclei of an atom. The direction of the bond determines the spatial structure of the molecules, i.e., their geometric shape. Hybridization - it is a change in the shape of some orbitals when forming a covalent bond to achieve more efficient orbital overlap. The chemical bond formed with the participation of electrons of hybrid orbitals is stronger than the bond with the participation of electrons of non-hybrid s- and p-orbitals, since more overlap occurs. Distinguish the following types hybridization (Fig. 10, Table 31): |
sp hybridization - one s-orbital and one p-orbital turn into two identical “hybrid” orbitals, the angle between their axes is 180°. The molecules in which sp-hybridization occurs have a linear geometry (BeCl 2).sp 2 hybridization- one s-orbital and two p-orbitals turn into three identical “hybrid” orbitals, the angle between their axes is 120°. Molecules in which sp 2 hybridization occurs have a flat geometry (BF 3, AlCl 3).
sp 3-hybridization- one s-orbital and three p-orbitals transform into four identical “hybrid” orbitals, the angle between the axes of which is 109°28". Molecules in which sp 3 hybridization occurs have a tetrahedral geometry (CH 4 , NH 3).
Rice. 10. Types of hybridization of valence orbitals: a - sp-hybridization of valence orbitals; b - sp 2 - hybridization of valence orbitals; V - sp 3-hybridization of valence orbitals
Chemicals are what the world around us is made of.
The properties of each chemical substance are divided into two types: chemical, which characterize its ability to form other substances, and physical, which are objectively observed and can be considered in isolation from chemical transformations. For example, the physical properties of a substance are its state of aggregation(solid, liquid or gaseous), thermal conductivity, heat capacity, solubility in various media (water, alcohol, etc.), density, color, taste, etc.
Transformations of some chemical substances in other substances are called chemical phenomena or chemical reactions. It should be noted that there are also physical phenomena that are obviously accompanied by changes in some physical properties substances without being converted into other substances. TO physical phenomena, for example, include the melting of ice, freezing or evaporation of water, etc.
The fact that a chemical phenomenon occurs during any process can be concluded by observing characteristic features chemical reactions, such as color change, sedimentation, gas evolution, heat and/or light.
For example, a conclusion about the occurrence of chemical reactions can be made by observing:
Formation of sediment when boiling water, called scale in everyday life;
The release of heat and light when a fire burns;
Change in color of a cut of a fresh apple in air;
Formation of gas bubbles during dough fermentation, etc.
The smallest particles of a substance that undergo virtually no changes during chemical reactions, but only connect with each other in a new way, are called atoms.
The very idea of the existence of such units of matter arose back in ancient Greece in the minds of ancient philosophers, which actually explains the origin of the term “atom”, since “atomos” literally translated from Greek means “indivisible”.
However, contrary to the idea of ancient Greek philosophers, atoms are not the absolute minimum of matter, i.e. they themselves have a complex structure.
Each atom consists of so-called subatomic particles - protons, neutrons and electrons, designated respectively by the symbols p +, n o and e -. The superscript in the notation used indicates that the proton has a unit positive charge, the electron has a unit negative charge, and the neutron has no charge.
As for the qualitative structure of an atom, in each atom all protons and neutrons are concentrated in the so-called nucleus, around which the electrons form an electron shell.
The proton and neutron have almost the same masses, i.e. m p ≈ m n, and the mass of the electron is almost 2000 times less than the mass of each of them, i.e. m p /m e ≈ m n /m e ≈ 2000.
Since the fundamental property of an atom is its electrical neutrality, and the charge of one electron is equal to the charge of one proton, from this we can conclude that the number of electrons in any atom is equal to the number of protons.
For example, the table below shows the possible composition of atoms:
Type of atoms with the same nuclear charge, i.e. With the same number protons in their nuclei are called a chemical element. Thus, from the table above we can conclude that atom1 and atom2 belong to one chemical element, and atom3 and atom4 belong to another chemical element.
Each chemical element has its own name and individual symbol, which is read in a certain way. So, for example, the simplest chemical element, the atoms of which contain only one proton in the nucleus, is called “hydrogen” and is denoted by the symbol “H”, which is read as “ash”, and a chemical element with a nuclear charge of +7 (i.e. containing 7 protons) - “nitrogen”, has the symbol “N”, which is read as “en”.
As you can see from the table above, atoms of one chemical element can differ in the number of neutrons in their nuclei.
Atoms belonging to the same chemical element, but having different quantities neutrons and, as a consequence, mass are called isotopes.
For example, the chemical element hydrogen has three isotopes - 1 H, 2 H and 3 H. The indices 1, 2 and 3 above the symbol H mean the total number of neutrons and protons. Those. Knowing that hydrogen is a chemical element, which is characterized by the fact that there is one proton in the nuclei of its atoms, we can conclude that in the 1 H isotope there are no neutrons at all (1-1 = 0), in the 2 H isotope - 1 neutron (2-1=1) and in the 3 H isotope – two neutrons (3-1=2). Since, as already mentioned, the neutron and proton have the same masses, and the mass of the electron is negligibly small in comparison with them, this means that the 2 H isotope is almost twice as heavy as the 1 H isotope, and the 3 H isotope is even three times heavier . Due to such a large scatter in the masses of hydrogen isotopes, the isotopes 2 H and 3 H were even assigned separate individual names and symbols, which is not typical for any other chemical element. The 2H isotope was named deuterium and given the symbol D, and the 3H isotope was given the name tritium and given the symbol T.
If we take the mass of the proton and neutron as one, and neglect the mass of the electron, in fact the upper left index, in addition to the total number of protons and neutrons in the atom, can be considered its mass, and therefore this index is called mass number and are designated by the symbol A. Since protons are responsible for the charge of the nucleus of any atom, and the charge of each proton is conventionally considered equal to +1, the number of protons in the nucleus is called the charge number (Z). By denoting the number of neutrons in an atom as N, the relationship between mass number, charge number and number of neutrons can be expressed mathematically as:
According to modern concepts, the electron has a dual (particle-wave) nature. It has the properties of both a particle and a wave. Like a particle, an electron has mass and charge, but at the same time, the flow of electrons, like a wave, is characterized by the ability to diffraction.
To describe the state of an electron in an atom, the concepts of quantum mechanics are used, according to which the electron does not have a specific trajectory of motion and can be located at any point in space, but with different probabilities.
The region of space around the nucleus where an electron is most likely to be found is called an atomic orbital.
An atomic orbital can have various shapes, size and orientation. An atomic orbital is also called an electron cloud.
Graphically, one atomic orbital is usually denoted as a square cell:
Quantum mechanics has an extremely complex mathematical apparatus, therefore, in the framework of a school chemistry course, only the consequences of quantum mechanical theory are considered.
According to these consequences, any atomic orbital and the electron located in it are completely characterized by 4 quantum numbers.
- The principal quantum number, n, determines the total energy of an electron in a given orbital. Range of values of the main quantum number – all integers, i.e. n = 1,2,3,4, 5, etc.
- The orbital quantum number - l - characterizes the shape of the atomic orbital and can take any integer value from 0 to n-1, where n, recall, is the main quantum number.
Orbitals with l = 0 are called s-orbitals. s-Orbitals are spherical in shape and have no directionality in space:
Orbitals with l = 1 are called p-orbitals. These orbitals have the shape of a three-dimensional figure eight, i.e. a shape obtained by rotating a figure eight around an axis of symmetry, and outwardly resemble a dumbbell:
Orbitals with l = 2 are called d-orbitals, and with l = 3 – f-orbitals. Their structure is much more complex.
3) Magnetic quantum number – m l – determines the spatial orientation of a specific atomic orbital and expresses the projection of the orbital angular momentum onto the direction magnetic field. The magnetic quantum number m l corresponds to the orientation of the orbital relative to the direction of the external magnetic field strength vector and can take any integer values from –l to +l, including 0, i.e. total possible values equals (2l+1). So, for example, for l = 0 m l = 0 (one value), for l = 1 m l = -1, 0, +1 (three values), for l = 2 m l = -2, -1, 0, +1 , +2 (five values of magnetic quantum number), etc.
So, for example, p-orbitals, i.e. orbitals with an orbital quantum number l = 1, having the shape of a “three-dimensional figure of eight,” correspond to three values of the magnetic quantum number (-1, 0, +1), which, in turn, correspond to three directions perpendicular to each other in space.
4) The spin quantum number (or simply spin) - m s - can conventionally be considered responsible for the direction of rotation of the electron in the atom; it can take on values. Electrons with different spins are indicated by vertical arrows directed in different directions: ↓ and .
The set of all orbitals in an atom that have the same principal quantum number is called the energy level or electron shell. Any arbitrary energy level with some number n consists of n 2 orbitals.
A set of orbitals with the same values of the principal quantum number and orbital quantum number represents an energy sublevel.
Each energy level, which corresponds to the principal quantum number n, contains n sublevels. In turn, each energy sublevel with orbital quantum number l consists of (2l+1) orbitals. Thus, the s sublevel consists of one s orbital, the p sublevel consists of three p orbitals, the d sublevel consists of five d orbitals, and the f sublevel consists of seven f orbitals. Since, as already mentioned, one atomic orbital is often denoted by one square cell, the s-, p-, d- and f-sublevels can be graphically represented as follows:
Each orbital corresponds to an individual strictly defined set of three quantum numbers n, l and m l.
The distribution of electrons among orbitals is called the electron configuration.
The filling of atomic orbitals with electrons occurs in accordance with three conditions:
- Minimum energy principle: Electrons fill orbitals starting from the lowest energy sublevel. The sequence of sublevels in increasing order of their energies is as follows: 1s<2s<2p<3s<3p<4s≤3d<4p<5s≤4d<5p<6s…;
To make it easier to remember this sequence of filling out electronic sublevels, the following graphic illustration is very convenient:
- Pauli principle: Each orbital can contain no more than two electrons.
If there is one electron in an orbital, then it is called unpaired, and if there are two, then they are called an electron pair.
- Hund's rule: the most stable state of an atom is one in which, within one sublevel, the atom has the maximum possible number of unpaired electrons. This most stable state of the atom is called the ground state.
In fact, the above means that, for example, the placement of 1st, 2nd, 3rd and 4th electrons in three orbitals of the p-sublevel will be carried out as follows:
The filling of atomic orbitals from hydrogen, which has a charge number of 1, to krypton (Kr), with a charge number of 36, will be carried out as follows:
Such a representation of the order of filling of atomic orbitals is called an energy diagram. Based on the electronic diagrams of individual elements, it is possible to write down their so-called electronic formulas (configurations). So, for example, an element with 15 protons and, as a consequence, 15 electrons, i.e. phosphorus (P) will have the following energy diagram:
When converted into an electronic formula, the phosphorus atom will take the form:
15 P = 1s 2 2s 2 2p 6 3s 2 3p 3
The normal size numbers to the left of the sublevel symbol show the energy level number, and the superscripts to the right of the sublevel symbol show the number of electrons in the corresponding sublevel.
Below are the electronic formulas of the first 36 elements of the periodic table by D.I. Mendeleev.
| period | Item no. | symbol | Name | electronic formula |
| I | 1 | H | hydrogen | 1s 1 |
| 2 | He | helium | 1s 2 | |
| II | 3 | Li | lithium | 1s 2 2s 1 |
| 4 | Be | beryllium | 1s 2 2s 2 | |
| 5 | B | boron | 1s 2 2s 2 2p 1 | |
| 6 | C | carbon | 1s 2 2s 2 2p 2 | |
| 7 | N | nitrogen | 1s 2 2s 2 2p 3 | |
| 8 | O | oxygen | 1s 2 2s 2 2p 4 | |
| 9 | F | fluorine | 1s 2 2s 2 2p 5 | |
| 10 | Ne | neon | 1s 2 2s 2 2p 6 | |
| III | 11 | Na | sodium | 1s 2 2s 2 2p 6 3s 1 |
| 12 | Mg | magnesium | 1s 2 2s 2 2p 6 3s 2 | |
| 13 | Al | aluminum | 1s 2 2s 2 2p 6 3s 2 3p 1 | |
| 14 | Si | silicon | 1s 2 2s 2 2p 6 3s 2 3p 2 | |
| 15 | P | phosphorus | 1s 2 2s 2 2p 6 3s 2 3p 3 | |
| 16 | S | sulfur | 1s 2 2s 2 2p 6 3s 2 3p 4 | |
| 17 | Cl | chlorine | 1s 2 2s 2 2p 6 3s 2 3p 5 | |
| 18 | Ar | argon | 1s 2 2s 2 2p 6 3s 2 3p 6 | |
| IV | 19 | K | potassium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 |
| 20 | Ca | calcium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 | |
| 21 | Sc | scandium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 | |
| 22 | Ti | titanium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 | |
| 23 | V | vanadium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 | |
| 24 | Cr | chromium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 here we observe the jump of one electron with s on d sublevel | |
| 25 | Mn | manganese | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 | |
| 26 | Fe | iron | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 6 | |
| 27 | Co | cobalt | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 7 | |
| 28 | Ni | nickel | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 | |
| 29 | Cu | copper | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 here we observe the jump of one electron with s on d sublevel | |
| 30 | Zn | zinc | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 | |
| 31 | Ga | gallium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 | |
| 32 | Ge | germanium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 2 | |
| 33 | As | arsenic | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 | |
| 34 | Se | selenium | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 4 | |
| 35 | Br | bromine | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 | |
| 36 | Kr | krypton | 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 |
As already mentioned, in their ground state, electrons in atomic orbitals are located according to the principle of least energy. However, in the presence of empty p-orbitals in the ground state of the atom, often, by imparting excess energy to it, the atom can be transferred to the so-called excited state. For example, a boron atom in its ground state has an electronic configuration and an energy diagram of the following form:
5 B = 1s 2 2s 2 2p 1
And in an excited state (*), i.e. When some energy is imparted to a boron atom, its electron configuration and energy diagram will look like this:
5 B* = 1s 2 2s 1 2p 2
Depending on which sublevel in the atom is filled last, chemical elements are divided into s, p, d or f.
Finding s, p, d and f elements in the table D.I. Mendeleev:
- The s-elements have the last s-sublevel to be filled. These elements include elements of the main (on the left in the table cell) subgroups of groups I and II.
- For p-elements, the p-sublevel is filled. The p-elements include the last six elements of each period, except the first and seventh, as well as elements of the main subgroups of groups III-VIII.
- d-elements are located between s- and p-elements in large periods.
- f-Elements are called lanthanides and actinides. They are listed at the bottom of the D.I. table. Mendeleev.
When writing electronic formulas for atoms of elements, indicate energy levels (values of the main quantum number n in the form of numbers - 1, 2, 3, etc.), energy sublevels (orbital quantum number values l in the form of letters - s, p, d, f) and the number at the top indicate the number of electrons in a given sublevel.
The first element in the table is D.I. Mendeleev is hydrogen, therefore the charge of the nucleus of the atom N equals 1, an atom has only one electron per s-sublevel of the first level. Therefore, the electronic formula of the hydrogen atom has the form:
The second element is helium; its atom has two electrons, so the electronic formula of the helium atom is 2 Not 1s 2. The first period includes only two elements, since the first energy level is filled with electrons, which can only be occupied by 2 electrons.
The third element in order - lithium - is already in the second period, therefore, its second energy level begins to be filled with electrons (we talked about this above). The filling of the second level with electrons begins with s-sublevel, therefore the electronic formula of the lithium atom is 3 Li 1s 2 2s 1 . The beryllium atom is completed filling with electrons s-sublevel: 4 Ve 1s 2 2s 2 .
In subsequent elements of the 2nd period, the second energy level continues to be filled with electrons, only now it is filled with electrons R-sublevel: 5 IN 1s 2 2s 2 2R 1 ; 6 WITH 1s 2 2s 2 2R 2 … 10 Ne 1s 2 2s 2 2R 6 .
The neon atom completes filling with electrons R-sublevel, this element ends the second period, it has eight electrons, since s- And R-sublevels can only contain eight electrons.
The elements of the 3rd period have a similar sequence of filling the energy sublevels of the third level with electrons. The electronic formulas of the atoms of some elements of this period are as follows:
11 Na 1s 2 2s 2 2R 6 3s 1 ; 12 Mg 1s 2 2s 2 2R 6 3s 2 ; 13 Al 1s 2 2s 2 2R 6 3s 2 3p 1 ;
14 Si 1s 2 2s 2 2R 6 3s 2 3p 2 ;…; 18 Ar 1s 2 2s 2 2R 6 3s 2 3p 6 .
The third period, like the second, ends with an element (argon), which is completely filled with electrons R-sublevel, although the third level includes three sublevels ( s, R, d). According to the above order of filling energy sublevels in accordance with Klechkovsky's rules, the energy of sublevel 3 d more sublevel 4 energy s, therefore, the potassium atom next to argon and the calcium atom behind it are filled with electrons 3 s– sublevel of the fourth level:
19 TO 1s 2 2s 2 2R 6 3s 2 3p 6 4s 1 ; 20 Sa 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 .
Starting from the 21st element - scandium, sublevel 3 in the atoms of the elements begins to be filled with electrons d. The electronic formulas of the atoms of these elements are:
21 Sc 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 3d 1 ; 22 Ti 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 3d 2 .
In the atoms of the 24th element (chromium) and the 29th element (copper), a phenomenon called “leakage” or “failure” of an electron is observed: an electron from the outer 4 s– sublevel “falls” by 3 d– sublevel, completing filling it halfway (for chromium) or completely (for copper), which contributes to greater stability of the atom:
24 Cr 1s 2 2s 2 2R 6 3s 2 3p 6 4s 1 3d 5 (instead of...4 s 2 3d 4) and
29 Cu 1s 2 2s 2 2R 6 3s 2 3p 6 4s 1 3d 10 (instead of...4 s 2 3d 9).
Starting from the 31st element - gallium, the filling of the 4th level with electrons continues, now - R– sublevel:
31 Ga 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 3d 10 4p 1 …; 36 Kr 1s 2 2s 2 2R 6 3s 2 3p 6 4s 2 3d 10 4p 6 .
This element ends the fourth period, which already includes 18 elements.
A similar order of filling energy sublevels with electrons occurs in the atoms of elements of the 5th period. For the first two (rubidium and strontium) it is filled s– sublevel of the 5th level, for the next ten elements (from yttrium to cadmium) is filled d– sublevel of the 4th level; The period is completed by six elements (from indium to xenon), the atoms of which are filled with electrons R– sublevel of the external, fifth level. There are also 18 elements in a period.
For elements of the sixth period, this order of filling is violated. At the beginning of the period, as usual, there are two elements whose atoms are filled with electrons s– sublevel of the external, sixth, level. The next element behind them, lanthanum, begins to fill with electrons d– sublevel of the previous level, i.e. 5 d. This completes the filling with electrons 5 d-sublevel stops and the next 14 elements - from cerium to lutetium - begin to fill f-sublevel of the 4th level. These elements are all included in one cell of the table, and below is an expanded row of these elements, called lanthanides.
Starting from the 72nd element - hafnium - to the 80th element - mercury, filling with electrons continues 5 d-sublevel, and the period ends, as usual, with six elements (from thallium to radon), the atoms of which are filled with electrons R– sublevel of the external, sixth, level. This is the largest period, including 32 elements.
In the atoms of the elements of the seventh, incomplete, period, the same order of filling sublevels is visible as described above. We let students write the electronic formulas of atoms of elements of the 5th – 7th periods themselves, taking into account everything said above.
Note:In some textbooks, a different order of writing the electronic formulas of atoms of elements is allowed: not in the order in which they are filled in, but in accordance with the number of electrons at each energy level given in the table. For example, the electronic formula of the arsenic atom may look like: As 1s 2 2s 2 2R 6 3s 2 3p 6 3d 10 4s 2 4p 3 .
The structure of the electronic shells of atoms of elements of the first four periods: $s-$, $p-$ and $d-$elements. Electronic configuration of an atom. Ground and excited states of atoms
The concept of atom arose in the ancient world to denote particles of matter. Translated from Greek, atom means “indivisible.”
Electrons
The Irish physicist Stoney, based on experiments, came to the conclusion that electricity is carried by the smallest particles existing in the atoms of all chemical elements. In $1891, Mr. Stoney proposed to call these particles electrons, which means "amber" in Greek.
A few years after the electron got its name, the English physicist Joseph Thomson and the French physicist Jean Perrin proved that electrons carry a negative charge. This is the smallest negative charge, which in chemistry is taken as a unit $(–1)$. Thomson even managed to determine the speed of the electron (it is equal to the speed of light - $300,000 km/s) and the mass of the electron (it is $1836$ times less than the mass of a hydrogen atom).
Thomson and Perrin connected the poles of a current source with two metal plates - a cathode and an anode, soldered into a glass tube from which the air was evacuated. When a voltage of about 10 thousand volts was applied to the electrode plates, a luminous discharge flashed in the tube, and particles flew from the cathode (negative pole) to the anode (positive pole), which scientists first called cathode rays, and then found out that it was a stream of electrons. Electrons hitting special substances, such as those on a TV screen, cause a glow.
The conclusion was drawn: electrons escape from the atoms of the material from which the cathode is made.
Free electrons or their flow can be obtained in other ways, for example, by heating a metal wire or by shining light on metals formed by elements of the main subgroup of group I of the periodic table (for example, cesium).
State of electrons in an atom
The state of an electron in an atom is understood as the totality of information about energy certain electron in space, in which it is located. We already know that an electron in an atom does not have a trajectory of motion, i.e. we can only talk about probabilities its location in the space around the nucleus. It can be located in any part of this space surrounding the nucleus, and the set of different positions is considered as an electron cloud with a certain negative charge density. Figuratively, this can be imagined this way: if it were possible to photograph the position of an electron in an atom after hundredths or millionths of a second, as in a photo finish, then the electron in such photographs would be represented as a point. If countless such photographs were superimposed, the picture would be of an electron cloud with the greatest density where there are the most of these points.
The figure shows a “cut” of such an electron density in a hydrogen atom passing through the nucleus, and the dashed line limits the sphere within which the probability of detecting an electron is $90%$. The contour closest to the nucleus covers a region of space in which the probability of detecting an electron is $10%$, the probability of detecting an electron inside the second contour from the nucleus is $20%$, inside the third is $≈30%$, etc. There is some uncertainty in the state of the electron. To characterize this special state, the German physicist W. Heisenberg introduced the concept of uncertainty principle, i.e. showed that it is impossible to simultaneously and accurately determine the energy and location of an electron. The more precisely the energy of an electron is determined, the more uncertain its position, and vice versa, having determined the position, it is impossible to determine the energy of the electron. The probability range for detecting an electron does not have clear boundaries. However, it is possible to select a space where the probability of finding an electron is maximum.
The space around the atomic nucleus in which an electron is most likely to be found is called an orbital.
It contains approximately $90%$ of the electron cloud, which means that about $90%$ of the time the electron is in this part of space. Based on their shape, there are four known types of orbitals, which are designated by the Latin letters $s, p, d$ and $f$. A graphical representation of some forms of electron orbitals is presented in the figure.
The most important characteristic of the motion of an electron in a certain orbital is the energy of its binding with the nucleus. Electrons with similar energy values form a single electron layer, or energy level. Energy levels are numbered starting from the nucleus: $1, 2, 3, 4, 5, 6$ and $7$.
The integer $n$ denoting the number of the energy level is called the principal quantum number.
It characterizes the energy of electrons occupying a given energy level. Electrons of the first energy level, closest to the nucleus, have the lowest energy. Compared to electrons of the first level, electrons of subsequent levels are characterized by a large amount of energy. Consequently, the electrons of the outer level are least tightly bound to the atomic nucleus.
The number of energy levels (electronic layers) in an atom is equal to the number of the period in the D.I. Mendeleev system to which the chemical element belongs: atoms of elements of the first period have one energy level; second period - two; seventh period - seven.
The largest number of electrons at an energy level is determined by the formula:
where $N$ is the maximum number of electrons; $n$ is the level number, or the main quantum number. Consequently: at the first energy level closest to the nucleus there can be no more than two electrons; on the second - no more than $8$; on the third - no more than $18$; on the fourth - no more than $32$. And how, in turn, are energy levels (electronic layers) arranged?
Starting from the second energy level $(n = 2)$, each of the levels is divided into sublevels (sublayers), slightly different from each other in the binding energy with the nucleus.
The number of sublevels is equal to the value of the main quantum number: the first energy level has one sub level; the second - two; third - three; fourth - four. Sublevels, in turn, are formed by orbitals.
Each value of $n$ corresponds to a number of orbitals equal to $n^2$. According to the data presented in the table, one can trace the connection between the principal quantum number $n$ and the number of sublevels, the type and number of orbitals, and the maximum number of electrons at the sublevel and level.
Main quantum number, types and number of orbitals, maximum number of electrons in sublevels and levels.
| Energy level $(n)$ | Number of sublevels equal to $n$ | Orbital type | Number of orbitals | Maximum number of electrons | ||
| in the sublevel | in level equal to $n^2$ | in the sublevel | at a level equal to $n^2$ | |||
| $K(n=1)$ | $1$ | $1s$ | $1$ | $1$ | $2$ | $2$ |
| $L(n=2)$ | $2$ | $2s$ | $1$ | $4$ | $2$ | $8$ |
| $2p$ | $3$ | $6$ | ||||
| $M(n=3)$ | $3$ | $3s$ | $1$ | $9$ | $2$ | $18$ |
| $3p$ | $3$ | $6$ | ||||
| $3d$ | $5$ | $10$ | ||||
| $N(n=4)$ | $4$ | $4s$ | $1$ | $16$ | $2$ | $32$ |
| $4p$ | $3$ | $6$ | ||||
| $4d$ | $5$ | $10$ | ||||
| $4f$ | $7$ | $14$ | ||||
Sublevels are usually denoted by Latin letters, as well as the shape of the orbitals of which they consist: $s, p, d, f$. So:
- $s$-sublevel - the first sublevel of each energy level closest to the atomic nucleus, consists of one $s$-orbital;
- $p$-sublevel - the second sublevel of each, except the first, energy level, consists of three $p$-orbitals;
- $d$-sublevel - the third sublevel of each, starting from the third, energy level, consists of five $d$-orbitals;
- The $f$-sublevel of each, starting from the fourth energy level, consists of seven $f$-orbitals.
Atomic nucleus
But not only electrons are part of atoms. Physicist Henri Becquerel discovered that a natural mineral containing a uranium salt also emits unknown radiation, exposing photographic films shielded from light. This phenomenon was called radioactivity.
There are three types of radioactive rays:
- $α$-rays, which consist of $α$-particles having a charge $2$ times greater than the charge of an electron, but with a positive sign, and a mass $4$ times greater than the mass of a hydrogen atom;
- $β$-rays represent a flow of electrons;
- $γ$-rays are electromagnetic waves with negligible mass that do not carry an electrical charge.
Consequently, the atom has a complex structure - it consists of a positively charged nucleus and electrons.
How is an atom structured?
In 1910, in Cambridge, near London, Ernest Rutherford and his students and colleagues studied the scattering of $α$ particles passing through thin gold foil and falling on a screen. Alpha particles usually deviated from the original direction by only one degree, seemingly confirming the uniformity and uniformity of the properties of gold atoms. And suddenly the researchers noticed that some $α$ particles abruptly changed the direction of their path, as if encountering some kind of obstacle.
By placing a screen in front of the foil, Rutherford was able to detect even those rare cases when $α$ particles, reflected from gold atoms, flew in the opposite direction.
Calculations showed that the observed phenomena could occur if the entire mass of the atom and all its positive charge were concentrated in a tiny central nucleus. The radius of the nucleus, as it turned out, is 100,000 times smaller than the radius of the entire atom, the region in which electrons with a negative charge are located. If we apply a figurative comparison, then the entire volume of an atom can be likened to the stadium in Luzhniki, and the nucleus can be likened to a soccer ball located in the center of the field.
An atom of any chemical element is comparable to a tiny solar system. Therefore, this model of the atom, proposed by Rutherford, is called planetary.
Protons and Neutrons
It turns out that the tiny atomic nucleus, in which the entire mass of the atom is concentrated, consists of two types of particles - protons and neutrons.
Protons have a charge equal to the charge of the electrons, but opposite in sign $(+1)$, and a mass equal to the mass of the hydrogen atom (it is taken as unity in chemistry). Protons are designated by the sign $↙(1)↖(1)p$ (or $p+$). Neutrons do not carry a charge, they are neutral and have a mass equal to the mass of a proton, i.e. $1$. Neutrons are designated by the sign $↙(0)↖(1)n$ (or $n^0$).
Protons and neutrons together are called nucleons(from lat. nucleus- core).
The sum of the number of protons and neutrons in an atom is called mass number. For example, the mass number of an aluminum atom is:
Since the mass of the electron, which is negligibly small, can be neglected, it is obvious that the entire mass of the atom is concentrated in the nucleus. Electrons are designated as follows: $e↖(-)$.
Since the atom is electrically neutral, it is also obvious that that the number of protons and electrons in an atom is the same. It is equal to the atomic number of the chemical element, assigned to it in the Periodic Table. For example, the nucleus of an iron atom contains $26$ protons, and $26$ electrons revolve around the nucleus. How to determine the number of neutrons?
As is known, the mass of an atom consists of the mass of protons and neutrons. Knowing the serial number of the element $(Z)$, i.e. the number of protons, and the mass number $(A)$, equal to the sum of the numbers of protons and neutrons, the number of neutrons $(N)$ can be found using the formula:
For example, the number of neutrons in an iron atom is:
$56 – 26 = 30$.
The table presents the main characteristics of elementary particles.
Basic characteristics of elementary particles.
Isotopes
Varieties of atoms of the same element that have the same nuclear charge but different mass numbers are called isotopes.
Word isotope consists of two Greek words: isos- identical and topos- place, means “occupying one place” (cell) in the Periodic Table of Elements.
Chemical elements found in nature are a mixture of isotopes. Thus, carbon has three isotopes with masses $12, 13, 14$; oxygen - three isotopes with masses $16, 17, 18, etc.
Usually, the relative atomic mass of a chemical element given in the Periodic Table is the average value of the atomic masses of a natural mixture of isotopes of a given element, taking into account their relative abundance in nature, therefore the values of atomic masses are quite often fractional. For example, natural chlorine atoms are a mixture of two isotopes - $35$ (there are $75%$ in nature) and $37$ (there are $25%$ in nature); therefore, the relative atomic mass of chlorine is $35.5$. Isotopes of chlorine are written as follows:
$↖(35)↙(17)(Cl)$ and $↖(37)↙(17)(Cl)$
The chemical properties of chlorine isotopes are exactly the same, as are the isotopes of most chemical elements, for example potassium, argon:
$↖(39)↙(19)(K)$ and $↖(40)↙(19)(K)$, $↖(39)↙(18)(Ar)$ and $↖(40)↙(18 )(Ar)$
However, hydrogen isotopes vary greatly in properties due to the dramatic multiple increase in their relative atomic mass; they were even given individual names and chemical symbols: protium - $↖(1)↙(1)(H)$; deuterium - $↖(2)↙(1)(H)$, or $↖(2)↙(1)(D)$; tritium - $↖(3)↙(1)(H)$, or $↖(3)↙(1)(T)$.
Now we can give a modern, more rigorous and scientific definition of a chemical element.
A chemical element is a collection of atoms with the same nuclear charge.
The structure of the electronic shells of atoms of elements of the first four periods
Let's consider the display of electronic configurations of atoms of elements according to the periods of the D.I. Mendeleev system.
Elements of the first period.
Diagrams of the electronic structure of atoms show the distribution of electrons across electronic layers (energy levels).
Electronic formulas of atoms show the distribution of electrons across energy levels and sublevels.
Graphic electronic formulas of atoms show the distribution of electrons not only across levels and sublevels, but also across orbitals.
In a helium atom, the first electron layer is complete - it contains $2$ electrons.
Hydrogen and helium are $s$ elements; the $s$ orbital of these atoms is filled with electrons.
Elements of the second period.
For all second-period elements, the first electron layer is filled, and electrons fill the $s-$ and $p$ orbitals of the second electron layer in accordance with the principle of least energy (first $s$ and then $p$) and the Pauli and Hund rules.
In the neon atom, the second electron layer is complete - it contains $8$ electrons.
Elements of the third period.
For atoms of elements of the third period, the first and second electron layers are completed, so the third electron layer is filled, in which electrons can occupy the 3s-, 3p- and 3d-sub levels.
The structure of the electronic shells of atoms of elements of the third period.
The magnesium atom completes its $3.5$ electron orbital. $Na$ and $Mg$ are $s$-elements.
In aluminum and subsequent elements, the $3d$ sublevel is filled with electrons.
| $↙(18)(Ar)$ Argon |  |
$1s^2(2)s^2(2)p^6(3)s^2(3)p^6$ |  |
An argon atom has $8$ electrons in its outer layer (third electron layer). As the outer layer is completed, but in total in the third electron layer, as you already know, there can be 18 electrons, which means that the elements of the third period have unfilled $3d$-orbitals.
All elements from $Al$ to $Ar$ are $р$ -elements.
$s-$ and $p$ -elements form main subgroups in the Periodic Table.
Elements of the fourth period.
Potassium and calcium atoms have a fourth electron layer and the $4s$ sublevel is filled, because it has lower energy than the $3d$ sublevel. To simplify the graphical electronic formulas of atoms of elements of the fourth period:
- Let us denote the conventional graphical electronic formula of argon as follows: $Ar$;
- We will not depict sublevels that are not filled in these atoms.
$K, Ca$ - $s$ -elements, included in the main subgroups. For atoms from $Sc$ to $Zn$, the 3d sublevel is filled with electrons. These are $3d$ elements. They are included in side subgroups, their outer electron layer is filled, they are classified as transitional elements.
Pay attention to the structure of the electronic shells of chromium and copper atoms. In them, one electron “fails” from the $4s-$ to the $3d$ sublevel, which is explained by the greater energy stability of the resulting $3d^5$ and $3d^(10)$ electronic configurations:
$↙(24)(Cr)$ $1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)3d^(4) 4s^(2)…$ 
$↙(29)(Cu)$ $1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)3d^(9)4s^(2)…$ 
| Element symbol, serial number, name | Electronic structure diagram | Electronic formula | Graphical electronic formula |
| $↙(19)(K)$ Potassium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^1$ | |
| $↙(20)(C)$ Calcium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2$ | |
| $↙(21)(Sc)$ Scandium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^1$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^1(4)s^1$ |  |
| $↙(22)(Ti)$ Titanium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^2$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^2(4)s^2$ |  |
| $↙(23)(V)$ Vanadium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^3$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^3(4)s^2$ |  |
| $↙(24)(Cr)$ Chrome |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^5$ or $1s^2(2)s^2(2)p ^6(3)p^6(3)d^5(4)s^1$ |  |
| $↙(29)(Cu)$ Chrome |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^1(3)d^(10)$ or $1s^2(2)s^2(2 )p^6(3)p^6(3)d^(10)(4)s^1$ |  |
| $↙(30)(Zn)$ Zinc |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^(10)$ or $1s^2(2)s^2(2 )p^6(3)p^6(3)d^(10)(4)s^2$ |  |
| $↙(31)(Ga)$ Gallium |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^(10)4p^(1)$ or $1s^2(2) s^2(2)p^6(3)p^6(3)d^(10)(4)s^(2)4p^(1)$ |  |
| $↙(36)(Kr)$ Krypton |  |
$1s^2(2)s^2(2)p^6(3)p^6(4)s^2(3)d^(10)4p^6$ or $1s^2(2)s^ 2(2)p^6(3)p^6(3)d^(10)(4)s^(2)4p^6$ |  |
In the zinc atom, the third electron layer is complete - all $3s, 3p$ and $3d$ sublevels are filled in it, with a total of $18$ electrons.
In the elements following zinc, the fourth electron layer, the $4p$ sublevel, continues to be filled. Elements from $Ga$ to $Kr$ - $р$ -elements.
The outer (fourth) layer of the krypton atom is complete and has $8$ electrons. But in total in the fourth electron layer, as you know, there can be $32$ electrons; the krypton atom still has unfilled $4d-$ and $4f$ sublevels.
For elements of the fifth period, sublevels are filled in in the following order: $5s → 4d → 5p$. And there are also exceptions associated with the “failure” of electrons in $↙(41)Nb$, $↙(42)Mo$, $↙(44)Ru$, $↙(45)Rh$, $↙(46) Pd$, $↙(47)Ag$. $f$ appears in the sixth and seventh periods -elements, i.e. elements for which the $4f-$ and $5f$ sublevels of the third outside electronic layer are filled, respectively.
$4f$ -elements called lanthanides.
$5f$ -elements called actinides.
The order of filling electronic sublevels in atoms of elements of the sixth period: $↙(55)Cs$ and $↙(56)Ba$ - $6s$ elements; $↙(57)La ... 6s^(2)5d^(1)$ - $5d$-element; $↙(58)Се$ – $↙(71)Lu - 4f$-elements; $↙(72)Hf$ – $↙(80)Hg - 5d$-elements; $↙(81)T1$ – $↙(86)Rn - 6d$-elements. But here, too, there are elements in which the order of filling of electronic orbitals is violated, which, for example, is associated with greater energy stability of half and completely filled $f$-sublevels, i.e. $nf^7$ and $nf^(14)$.
Depending on which sublevel of the atom is filled with electrons last, all elements, as you already understood, are divided into four electron families, or blocks:
- $s$ -elements; the $s$-sublevel of the outer level of the atom is filled with electrons; $s$-elements include hydrogen, helium and elements of the main subgroups of groups I and II;
- $p$ -elements; the $p$-sublevel of the outer level of the atom is filled with electrons; $p$-elements include elements of the main subgroups of groups III–VIII;
- $d$ -elements; the $d$-sublevel of the pre-external level of the atom is filled with electrons; $d$-elements include elements of secondary subgroups of groups I–VIII, i.e. elements of intercalary decades of large periods located between $s-$ and $p-$elements. They are also called transition elements;
- $f$ -elements; electrons fill the $f-$sublevel of the third outer level of the atom; these include lanthanides and actinides.
Electronic configuration of an atom. Ground and excited states of atoms
Swiss physicist W. Pauli in $1925 found that an atom can have no more than two electrons in one orbital, having opposite (antiparallel) backs (translated from English as a spindle), i.e. possessing properties that can be conventionally imagined as the rotation of an electron around its imaginary axis clockwise or counterclockwise. This principle is called Pauli principle.
If there is one electron in an orbital, it is called unpaired, if two, then this paired electrons, i.e. electrons with opposite spins.
The figure shows a diagram of dividing energy levels into sublevels.
$s-$ Orbital, as you already know, has a spherical shape. The electron of the hydrogen atom $(n = 1)$ is located in this orbital and is unpaired. For this reason it electronic formula, or electronic configuration, is written like this: $1s^1$. In electronic formulas, the number of the energy level is indicated by the number in front of the letter $(1...)$, the Latin letter denotes the sublevel (type of orbital), and the number written to the right above the letter (as an exponent) shows the number of electrons in the sublevel.
For a helium atom He, which has two paired electrons in one $s-$orbital, this formula is: $1s^2$. The electron shell of the helium atom is complete and very stable. Helium is a noble gas. At the second energy level $(n = 2)$ there are four orbitals, one $s$ and three $p$. Electrons of the $s$-orbital of the second level ($2s$-orbital) have higher energy, because are at a greater distance from the nucleus than the electrons of the $1s$ orbital $(n = 2)$. In general, for each value of $n$ there is one $s-$orbital, but with a corresponding supply of electron energy on it and, therefore, with a corresponding diameter, growing as the value of $n$ increases. The $s-$Orbital, as you already know , has a spherical shape. The electron of the hydrogen atom $(n = 1)$ is located in this orbital and is unpaired. Therefore, its electronic formula, or electronic configuration, is written as follows: $1s^1$. In electronic formulas, the number of the energy level is indicated by the number in front of the letter $(1...)$, the Latin letter denotes the sublevel (type of orbital), and the number written to the right above the letter (as an exponent) shows the number of electrons in the sublevel.
For a helium atom $He$, which has two paired electrons in one $s-$orbital, this formula is: $1s^2$. The electron shell of the helium atom is complete and very stable. Helium is a noble gas. At the second energy level $(n = 2)$ there are four orbitals, one $s$ and three $p$. Electrons of $s-$orbitals of the second level ($2s$-orbitals) have higher energy, because are at a greater distance from the nucleus than the electrons of the $1s$ orbital $(n = 2)$. In general, for each value of $n$ there is one $s-$orbital, but with a corresponding supply of electron energy on it and, therefore, with a corresponding diameter, growing as the value of $n$ increases. 
$p-$ Orbital has the shape of a dumbbell, or a voluminous figure eight. All three $p$-orbitals are located in the atom mutually perpendicular along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized once again that each energy level (electronic layer), starting from $n= 2$, has three $p$-orbitals. As the value of $n$ increases, electrons occupy $p$-orbitals located at large distances from the nucleus and directed along the $x, y, z$ axes.
For elements of the second period $(n = 2)$, first one $s$-orbital is filled, and then three $p$-orbitals; electronic formula $Li: 1s^(2)2s^(1)$. The $2s^1$ electron is more weakly bound to the nucleus of the atom, so the lithium atom can easily give it up (as you obviously remember, this process is called oxidation), turning into a lithium ion $Li^+$.
In the beryllium Be atom, the fourth electron is also located in the $2s$ orbital: $1s^(2)2s^(2)$. The two outer electrons of the beryllium atom are easily detached - $B^0$ is oxidized into the $Be^(2+)$ cation.
In the boron atom, the fifth electron occupies the $2p$ orbital: $1s^(2)2s^(2)2p^(1)$. Next, the $C, N, O, F$ atoms are filled with $2p$-orbitals, which ends with the noble gas neon: $1s^(2)2s^(2)2p^(6)$.
For elements of the third period, the $3s-$ and $3p$ orbitals are filled, respectively. Five $d$-orbitals of the third level remain free:
$↙(11)Na 1s^(2)2s^(2)2p^(6)3s^(1)$,
$↙(17)Cl 1s^(2)2s^(2)2p^(6)3s^(2)3p^(5)$,
$↙(18)Ar 1s^(2)2s^(2)2p^(6)3s^(2)3p^(6)$.
Sometimes in diagrams depicting the distribution of electrons in atoms, only the number of electrons at each energy level is indicated, i.e. write abbreviated electronic formulas of atoms of chemical elements, in contrast to the full electronic formulas given above, for example:
$↙(11)Na 2, 8, 1;$ $↙(17)Cl 2, 8, 7;$ $↙(18)Ar 2, 8, 8$.
For elements of large periods (fourth and fifth), the first two electrons occupy $4s-$ and $5s$ orbitals, respectively: $↙(19)K 2, 8, 8, 1;$ $↙(38)Sr 2, 8, 18, 8, 2$. Starting from the third element of each major period, the next ten electrons will go to the previous $3d-$ and $4d-$orbitals, respectively (for elements of side subgroups): $↙(23)V 2, 8, 11, 2;$ $↙( 26)Fr 2, 8, 14, 2;$ $↙(40)Zr 2, 8, 18, 10, 2;$ $↙(43)Tc 2, 8, 18, 13, 2$. As a rule, when the previous $d$-sublevel is filled, the outer ($4р-$ and $5р-$, respectively) $р-$sublevel will begin to be filled: $↙(33)As 2, 8, 18, 5;$ $ ↙(52)Te 2, 8, 18, 18, 6$.
For elements of large periods - the sixth and the incomplete seventh - electronic levels and sublevels are filled with electrons, as a rule, like this: the first two electrons enter the outer $s-$sublevel: $↙(56)Ba 2, 8, 18, 18, 8, 2;$ $↙(87)Fr 2, 8, 18, 32, 18, 8, 1$; the next one electron (for $La$ and $Ca$) to the previous $d$-sublevel: $↙(57)La 2, 8, 18, 18, 9, 2$ and $↙(89)Ac 2, 8, 18, 32, 18, 9, 2$.
Then the next $14$ electrons will go to the third outer energy level, to the $4f$ and $5f$ orbitals of lanthanides and actinides, respectively: $↙(64)Gd 2, 8, 18, 25, 9, 2;$ $↙(92 )U 2, 8, 18, 32, 21, 9, 2$.
Then the second external energy level ($d$-sublevel) of elements of side subgroups will begin to build up again: $↙(73)Ta 2, 8, 18, 32, 11, 2;$ $↙(104)Rf 2, 8, 18 , 32, 32, 10, 2$. And finally, only after the $d$-sublevel is completely filled with ten electrons will the $p$-sublevel be filled again: $↙(86)Rn 2, 8, 18, 32, 18, 8$.
Very often the structure of the electronic shells of atoms is depicted using energy or quantum cells - the so-called graphic electronic formulas. For this notation, the following notation is used: each quantum cell is designated by a cell that corresponds to one orbital; Each electron is indicated by an arrow corresponding to the spin direction. When writing a graphical electronic formula, you should remember two rules: Pauli principle, according to which there can be no more than two electrons in a cell (orbital), but with antiparallel spins, and F. Hund's rule, according to which electrons occupy free cells first one at a time and have the same spin value, and only then pair, but the spins, according to the Pauli principle, will be in opposite directions.
Algorithm for composing the electronic formula of an element:
1. Determine the number of electrons in an atom using the Periodic Table of Chemical Elements D.I. Mendeleev.
2. Using the number of the period in which the element is located, determine the number of energy levels; the number of electrons in the last electronic level corresponds to the group number.
3. Divide the levels into sublevels and orbitals and fill them with electrons in accordance with the rules for filling orbitals:
It must be remembered that the first level contains a maximum of 2 electrons 1s 2, on the second - a maximum of 8 (two s and six R: 2s 2 2p 6), on the third - a maximum of 18 (two s, six p, and ten d: 3s 2 3p 6 3d 10).
- Principal quantum number n should be minimal.
- First to fill s- sublevel, then р-, d- b f- sublevels.
- Electrons fill the orbitals in order of increasing energy of the orbitals (Klechkovsky's rule).
- Within a sublevel, electrons first occupy free orbitals one by one, and only after that they form pairs (Hund’s rule).
- There cannot be more than two electrons in one orbital (Pauli principle).
Examples.
1. Let's create the electronic formula of nitrogen. Nitrogen is number 7 on the periodic table.
2. Let's create the electronic formula for argon. Argon is number 18 on the periodic table.
1s 2 2s 2 2p 6 3s 2 3p 6.
3. Let's create the electronic formula of chromium. Chromium is number 24 on the periodic table.
1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5
Energy diagram of zinc.
4. Let's create the electronic formula of zinc. Zinc is number 30 on the periodic table.
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10
Please note that part of the electronic formula, namely 1s 2 2s 2 2p 6 3s 2 3p 6, is the electronic formula of argon.
The electronic formula of zinc can be represented as: