Let's find out how to write the electronic formula of a chemical element. This question is important and relevant, since it gives an idea not only about the structure, but also about the alleged physical and chemical properties of the atom in question.

Compilation rules

In order to compose a graphical and electronic formula of a chemical element, it is necessary to have an idea of ​​​​the theory of the structure of the atom. To begin with, there are two main components of an atom: the nucleus and the negative electrons. The nucleus includes neutrons, which have no charge, as well as protons, which have a positive charge.

Arguing how to compose and determine the electronic formula of a chemical element, we note that in order to find the number of protons in the nucleus, the periodic system of Mendeleev is required.

The number of an element in order corresponds to the number of protons in its nucleus. The number of the period in which the atom is located characterizes the number of energy layers on which the electrons are located.

To determine the number of neutrons devoid of an electric charge, it is necessary to subtract its serial number (the number of protons) from the relative mass of an atom of an element.

Instruction

In order to understand how to compose the electronic formula of a chemical element, consider the rule for filling sublevels with negative particles, formulated by Klechkovsky.

Depending on the amount of free energy the free orbitals have, a series is drawn up that characterizes the sequence of filling the levels with electrons.

Each orbital contains only two electrons, which are arranged in antiparallel spins.

In order to express the structure of electron shells, graphic formulas are used. What do the electronic formulas of atoms of chemical elements look like? How to make graphic options? These questions are included in the school chemistry course, so we will dwell on them in more detail.

There is a certain matrix (basis) that is used when compiling graphic formulas. The s-orbital is characterized by only one quantum cell, in which two electrons are located opposite to each other. They are indicated graphically by arrows. For the p orbital, three cells are depicted, each also contains two electrons, ten electrons are located on the d orbital, and f is filled with fourteen electrons.

Examples of compiling electronic formulas

Let's continue the conversation about how to compose the electronic formula of a chemical element. For example, you need to make a graphical and electronic formula for the element manganese. First, we determine the position of this element in the periodic system. It has atomic number 25, so there are 25 electrons in an atom. Manganese is an element of the fourth period, therefore, it has four energy levels.

How to write the electronic formula of a chemical element? We write down the sign of the element, as well as its ordinal number. Using the Klechkovsky rule, we distribute electrons over energy levels and sublevels. We sequentially arrange them on the first, second, and third level, inscribing two electrons in each cell.

Then we sum them up, getting 20 pieces. Three levels are completely filled with electrons, and only five electrons remain on the fourth. Considering that each type of orbital has its own energy reserve, we distribute the remaining electrons to the 4s and 3d sublevels. As a result, the finished electron-graphic formula for the manganese atom has the following form:

1s2/2s2, 2p6/3s2, 3p6/4s2, 3d3

Practical value

With the help of electron-graphic formulas, you can clearly see the number of free (unpaired) electrons that determine the valency of a given chemical element.

We offer a generalized algorithm of actions, with the help of which you can compose electronic graphic formulas of any atoms located in the periodic table.

The first step is to determine the number of electrons using the periodic table. The period number indicates the number of energy levels.

Belonging to a certain group is associated with the number of electrons that are in the outer energy level. The levels are subdivided into sublevels, filled in according to the Klechkovsky rule.

Conclusion

In order to determine the valence capabilities of any chemical element located in the periodic table, it is necessary to draw up an electron-graphic formula of its atom. The algorithm given above will allow to cope with the task, to determine the possible chemical and physical properties of the atom.

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; the numbers in front of the letters indicate the energy level in which the given electron is located, and the index at the top right is the number of electrons in this 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 system and fulfill the basic provisions 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 an arrangement of electrons in energy cells.

For iron atoms, such a scheme has the following form:

This diagram clearly shows the implementation of Hund's rule. At the 3d sublevel, the maximum number of cells (four) is filled with unpaired electrons. The image of the structure of the electron shell in the atom in the form of electronic formulas and in the form of diagrams does not clearly reflect the wave properties of the electron.

The wording of the periodic law as amended YES. Mendeleev : the properties of simple bodies, as well as the forms and properties of the 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 the elements, as well as the forms and properties of their compounds, are in a periodic dependence on the magnitude of the charge of the nucleus of their atoms.

Thus, the positive charge of the nucleus (rather than atomic mass) turned out to be a more accurate argument on which the properties of elements and their compounds depend.

Valence- is the number of chemical bonds that one atom is bonded to another.
The valence possibilities 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 determines mainly the properties of their atoms. Therefore, these levels are called valence. The electrons of these levels, and sometimes of the pre-external levels, can take part in the formation of chemical bonds. Such electrons are also called valence electrons.

Stoichiometric valence chemical element - is the number of equivalents that a given atom can attach to itself, or is the number of equivalents in an atom.

Equivalents are determined by the number of attached or substituted hydrogen atoms, therefore, the stoichiometric valence is equal to the number of hydrogen atoms with which this atom interacts. But not all elements interact freely, but almost everything interacts with oxygen, so the stoichiometric valency can be defined as twice the number of attached oxygen atoms.

For example, the stoichiometric valency of sulfur in hydrogen sulfide H 2 S is 2, in oxide SO 2 - 4, in oxide SO 3 -6.

When determining the stoichiometric valency of an element according to the formula of a binary compound, one should be guided by the rule: the total valency 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 valence with a plus sign (for a metal or a more electropositive element in a 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 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 of the main subgroup of the second group of the Periodic system (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 for 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 system, 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 -3 (in ammonia and ammonium salts).

7. The oxidation states of the elements in the compound compensate each other so that their sum for all atoms in a molecule or a neutral formula unit is zero, and for an ion - 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 formulate multi-element compounds.

Degree of oxidation (oxidation number,) — auxiliary conditional 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 of the electric charge attributed to the atom, assuming that the electron pairs that carry out the bond are completely biased towards more electronegative atoms (that is, based on the assumption that the compound consists only of ions).

The oxidation state corresponds to the number of electrons that must be added to a positive ion to reduce it to a neutral atom, or taken 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 the elements, depending on the structure of the electron shell of the atom, change according to the periods and groups of the periodic system. Since electronic structures in a number of analogous elements are only similar, but not identical, then when moving from one element in a group to another, not a simple repetition of properties is observed for them, but their more or less clearly expressed regular 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 affinity.

Ionization energy (Ei) is the minimum amount of energy required for the detachment and complete removal of an electron from an atom in the gas phase at T = 0

K without transferring kinetic energy to the released electron with the transformation of the atom into a positively charged ion: E + Ei = E + + e-. The ionization energy is a positive value and has the lowest values ​​for alkali metal atoms and the highest for noble (inert) gas atoms.

Electron affinity (Ee) is the energy released or absorbed when an electron is attached to an atom in the gas phase at T = 0

K with the transformation of the atom into a negatively charged ion without transferring kinetic energy to the particle:

E + e- = E- + Ee.

Halogens, especially fluorine, have the maximum electron affinity (Ee = -328 kJ/mol).

The values ​​of Ei and Ee are expressed in kilojoules per mol (kJ/mol) or in electron volts per atom (eV).

The ability of a bound atom to displace the 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 attach 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 the electron affinity of free atoms h = (Ee + Ei)/2

In periods, there is a general tendency for an increase in the ionization energy and electronegativity with an increase in the charge of the atomic nucleus; in groups, these values ​​decrease with an increase in the ordinal number of the element.

It should be emphasized that an element cannot be assigned a constant value of electronegativity, since it depends on many factors, in particular, on the valence state of the element, the type of compound in which it enters, the number and type of neighboring atoms.

Atomic and ionic radii. The dimensions of atoms and ions are determined by the dimensions of the electron shell. According to quantum mechanical concepts, the electron shell does not have strictly defined boundaries. Therefore, for the radius of a free atom or ion, we can take theoretically calculated distance from the core to the position of the main maximum density of the outer electron clouds. This distance is called the orbital radius. In practice, the values ​​of the radii of atoms and ions in compounds, calculated from experimental data, are usually used. 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 atom of an element and is periodic. In periods, as the atomic number increases, the radii tend to decrease. The greatest decrease is typical for elements of small periods, since the outer electronic level is filled in them. In large periods in the families of d- and f-elements, this change is less sharp, since the filling of electrons in them occurs in the preexternal layer. In subgroups, the radii of atoms and ions of the same type generally increase.

The periodic system of elements is a clear example of the manifestation 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 is preserved.

In the period from left to right (→), the oxidizing and non-metallic properties of the elements increase, while 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 interconnection of atoms in a molecule, or crystal lattice, as a result of the action of electric forces of attraction between 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 bonding is carried out by valence electrons. According to modern concepts, the chemical bond has an electronic nature, but it is carried out in different ways. Therefore, there are three main types of chemical bonds: covalent, ionic, metallic. Between molecules arises hydrogen bond, and happen van der Waals interactions.

The main characteristics of a chemical bond are:

- bond length - is the internuclear distance between chemically bonded atoms.

It depends on the nature of the interacting atoms and on the multiplicity of the bond. With an increase in the multiplicity, the bond length decreases, and, consequently, its strength increases;

- bond multiplicity - is determined by the number of electron pairs linking 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;

Binding energy E CB - this is the energy that is released during the formation of this bond and is spent on breaking it, kJ / mol.

covalent bond - A chemical bond formed by sharing a pair of electrons with two atoms.

The explanation of the chemical bond by the appearance of common electron pairs between atoms formed the basis of the spin theory of valence, the tool of which is valence bond method (MVS) , discovered by Lewis in 1916. For the quantum mechanical description of the chemical bond and the structure of molecules, another method is used - molecular orbital method (MMO) .

Valence bond method

The basic principles of the formation of a chemical bond according to MVS:

1. A chemical bond is formed due to 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 the 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 a covalent bond:

exchange mechanism. The bond is formed by sharing the valence electrons of two neutral atoms. Each atom gives one unpaired electron to a common electron pair:

Rice. 7. Exchange mechanism for the formation of a covalent bond: a- non-polar; b- polar

Donor-acceptor mechanism. One atom (donor) provides an electron pair, and another atom (acceptor) provides an empty orbital for this pair.

connections, educated according to the donor-acceptor mechanism, belong to complex compounds

Rice. 8. Donor-acceptor mechanism of covalent bond formation

A covalent bond has certain characteristics.

Saturability - the property of atoms to form a strictly defined number of covalent bonds. Due to the saturation of the bonds, the molecules have a certain composition.

Orientation - t . e. the connection is formed in the direction of maximum overlap of electron clouds . With respect to the line connecting the centers of atoms forming a bond, there are: 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 orientation 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 in the formation of a covalent bond in order to achieve a more efficient overlap of orbitals. 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 there is more overlap. There are the following types of hybridization (Fig. 10, Table 31): sp hybridization - one s-orbital and one p-orbital turn into two identical "hybrid" orbitals, the angle between the axes of which is 180°. 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 the axes of which is 120°. Molecules in which sp 2 hybridization is carried out have a flat geometry (BF 3 , AlCl 3).

sp 3-hybridization- one s-orbital and three p-orbitals turn 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 , NH3).

Rice. 10. Types of hybridizations of valence orbitals: a - sp-hybridization of valence orbitals; b - sp2- hybridization of valence orbitals; in - sp 3 - hybridization of valence orbitals

Electronic configuration an atom is a numerical representation of its electron orbitals. Electron orbitals are regions of various shapes located around the atomic nucleus, in which it is mathematically probable that an electron will be found. The electronic configuration helps to quickly and easily tell the reader how many electron orbitals an atom has, as well as determine the number of electrons in each orbital. After reading this article, you will master the method of compiling electronic configurations.

Steps

Distribution of electrons using the periodic system of D. I. Mendeleev

Find the atomic number of your atom. Each atom has a certain number of electrons associated with it. Find the symbol for your atom in the periodic table. The atomic number is a positive integer starting from 1 (for hydrogen) and increasing by one for each subsequent atom. The atomic number is the number of protons in an atom, and therefore it is also the number of electrons in an atom with zero charge.

Determine the charge of an atom. Neutral atoms will have the same number of electrons as shown in the periodic table. However, charged atoms will have more or fewer electrons, depending on the magnitude of their charge. If you are working with a charged atom, add or subtract electrons as follows: add one electron for every negative charge and subtract one for every positive charge.

• For example, a sodium atom with a charge of -1 will have an extra electron in addition to its base atomic number of 11. In other words, an atom will have 12 electrons in total.
• If we are talking about a sodium atom with a charge of +1, one electron must be subtracted from the base atomic number 11. So the atom will have 10 electrons.

1. Memorize the basic list of orbitals. As the number of electrons increases in an atom, they fill the various sublevels of the electron shell of the atom according to a certain sequence. Each sublevel of the electron shell, when filled, contains an even number of electrons. There are the following sublevels:

   Understand the electronic configuration record. Electronic configurations are written down in order to clearly reflect the number of electrons in each orbital. Orbitals are written sequentially, with the number of atoms in each orbital written as a superscript to the right of the orbital name. The completed electronic configuration has the form of a sequence of sublevel designations and superscripts.

   • Here, for example, is the simplest electronic configuration: 1s 2 2s 2 2p 6 . This configuration shows that there are two electrons in the 1s sublevel, two electrons in the 2s sublevel, and six electrons in the 2p sublevel. 2 + 2 + 6 = 10 electrons in total. This is the electronic configuration of the neutral neon atom (neon atomic number is 10).

2. Remember the order of the orbitals. Keep in mind that electron orbitals are numbered in ascending order of electron shell number, but arranged in ascending energy order. For example, a filled 4s 2 orbital has less energy (or less mobility) than a partially filled or filled 3d 10, so the 4s orbital is written first. Once you know the order of the orbitals, you can easily fill them in according to the number of electrons in the atom. The order in which the orbitals are filled is as follows: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

   • The electronic configuration of an atom in which all orbitals are filled will have the following form: 10 7p 6
   • Note that the above notation, when all orbits are filled, is the electron configuration of the element Uuo (ununoctium) 118, the highest numbered atom in the Periodic Table. Therefore, this electronic configuration contains all currently known electronic sublevels of a neutrally charged atom.

3. Fill in the orbitals according to the number of electrons in your atom. For example, if we want to write down the electronic configuration of a neutral calcium atom, we must start by looking up its atomic number in the periodic table. Its atomic number is 20, so we will write the configuration of an atom with 20 electrons according to the above order.

   • Fill in the orbitals in the above order until you reach the twentieth electron. The first 1s orbital will have two electrons, the 2s orbital will also have two, the 2p orbital will have six, the 3s orbital will have two, the 3p orbital will have 6, and the 4s orbital will have 2 (2 + 2 + 6 +2 +6 + 2 = 20 .) In other words, the electronic configuration of calcium has the form: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 .
   • Note that the orbitals are in ascending order of energy. For example, when you are ready to move to the 4th energy level, then first write down the 4s orbital, and then 3d. After the fourth energy level, you move on to the fifth, where the same order is repeated. This happens only after the third energy level.

4. Use the periodic table as a visual cue. You have probably already noticed that the shape of the periodic table corresponds to the order of electronic sublevels in electronic configurations. For example, atoms in the second column from the left always end in "s 2 ", while atoms on the right edge of the thin middle section always end in "d 10 ", and so on. Use the periodic table as a visual guide to writing configurations - as the order in which you add to the orbitals corresponds to your position in the table. See below:

   • In particular, the two leftmost columns contain atoms whose electronic configurations end in s orbitals, the right block of the table contains atoms whose configurations end in p orbitals, and at the bottom of the atoms end in f orbitals.
   • For example, when you write down the electronic configuration of chlorine, think like this: "This atom is located in the third row (or "period") of the periodic table. It is also located in the fifth group of the orbital block p of the periodic table. Therefore, its electronic configuration will end in. ..3p 5
   • Note that the elements in the d and f orbital regions of the table have energy levels that do not correspond to the period in which they are located. For example, the first row of a block of elements with d-orbitals corresponds to 3d orbitals, although it is located in the 4th period, and the first row of elements with f-orbitals corresponds to the 4f orbital, despite the fact that it is located in the 6th period.

5. Learn the abbreviations for writing long electronic configurations. The atoms on the right side of the periodic table are called noble gases. These elements are chemically very stable. To shorten the process of writing long electron configurations, simply write in square brackets the chemical symbol for the nearest noble gas with fewer electrons than your atom, and then continue to write the electronic configuration of subsequent orbital levels. See below:

   • To understand this concept, it will be helpful to write an example configuration. Let's write the configuration of zinc (atomic number 30) using the noble gas abbreviation. The full zinc configuration looks like this: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 . However, we see that 1s 2 2s 2 2p 6 3s 2 3p 6 is the electronic configuration of argon, a noble gas. Simply replace the electronic configuration part of zinc with the chemical symbol for argon in square brackets (.)
   • So, the electronic configuration of zinc, written in abbreviated form, is: 4s 2 3d 10 .
   • Note that if you are writing the electronic configuration of a noble gas, say argon, you cannot write! One must use the abbreviation of the noble gas in front of this element; for argon it will be neon ().

   Using ADOMAH Periodic Table

   1. Master the ADOMAH periodic table. This method of recording the electronic configuration does not require memorization, however, it requires a modified periodic table, since in the traditional periodic table, starting from the fourth period, the period number does not correspond to the electron shell. Find the ADOMAH periodic table, a special type of periodic table designed by scientist Valery Zimmerman. It is easy to find with a short internet search.

      • In the ADOMAH periodic table, the horizontal rows represent groups of elements such as halogens, noble gases, alkali metals, alkaline earth metals, etc. Vertical columns correspond to electronic levels, and so-called "cascades" (diagonal lines connecting blocks s, p, d and f) correspond to periods.
      • Helium is moved to hydrogen, since both of these elements are characterized by a 1s orbital. The period blocks (s,p,d and f) are shown on the right side and the level numbers are given at the bottom. Elements are represented in boxes numbered from 1 to 120. These numbers are the usual atomic numbers, which represent the total number of electrons in a neutral atom.

   2. Find your atom in the ADOMAH table. To write down the electronic configuration of an element, find its symbol in the ADOMAH periodic table and cross out all elements with a higher atomic number. For example, if you need to write down the electronic configuration of erbium (68), cross out all the elements from 69 to 120.

      • Pay attention to the numbers from 1 to 8 at the base of the table. These are the electronic level numbers, or column numbers. Ignore columns that contain only crossed out items. For erbium, columns with numbers 1,2,3,4,5 and 6 remain.

   3. Count the orbital sublevels up to your element. Looking at the block symbols shown to the right of the table (s, p, d, and f) and the column numbers shown at the bottom, ignore the diagonal lines between the blocks and break the columns into block-columns, listing them in order from bottom to top. And again, ignore the blocks in which all the elements are crossed out. Write the column blocks starting from the column number followed by the block symbol, thus: 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 6s (for erbium).

      • Please note: The above electronic configuration Er is written in ascending order of the electronic sublevel number. It can also be written in the order in which the orbitals are filled. To do this, follow the cascades from bottom to top, not columns, when you write column blocks: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 12 .

   4. Count the electrons for each electronic sublevel. Count the elements in each column block that have not been crossed out by attaching one electron from each element, and write their number next to the block symbol for each column block as follows: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 4f 12 5s 2 5p 6 6s 2 . In our example, this is the electronic configuration of erbium.

   5. Be aware of incorrect electronic configurations. There are eighteen typical exceptions related to the electronic configurations of atoms in the lowest energy state, also called the ground energy state. They do not obey the general rule only in the last two or three positions occupied by electrons. In this case, the actual electronic configuration assumes that the electrons are in a state of lower energy compared to the standard configuration of the atom. Exception atoms include:

      • Cr(..., 3d5, 4s1); Cu(..., 3d10, 4s1); Nb(..., 4d4, 5s1); Mo(..., 4d5, 5s1); Ru(..., 4d7, 5s1); Rh(..., 4d8, 5s1); Pd(..., 4d10, 5s0); Ag(..., 4d10, 5s1); La(..., 5d1, 6s2); Ce(..., 4f1, 5d1, 6s2); Gd(..., 4f7, 5d1, 6s2); Au(..., 5d10, 6s1); AC(..., 6d1, 7s2); Th(..., 6d2, 7s2); Pa(..., 5f2, 6d1, 7s2); U(..., 5f3, 6d1, 7s2); Np(..., 5f4, 6d1, 7s2) and cm(..., 5f7, 6d1, 7s2).
   • To find the atomic number of an atom when it is written in electronic form, simply add up all the numbers that follow the letters (s, p, d, and f). This only works for neutral atoms, if you are dealing with an ion, then nothing will work - you will have to add or subtract the number of extra or lost electrons.
   • The number following the letter is a superscript, do not make a mistake in the control.
   • The "stability of a half-filled" sublevel does not exist. This is a simplification. Any stability that pertains to "half-full" sublevels is due to the fact that each orbital is occupied by one electron, so repulsion between electrons is minimized.
   • Each atom tends to a stable state, and the most stable configurations have filled sublevels s and p (s2 and p6). Noble gases have this configuration, so they rarely react and are located on the right in the periodic table. Therefore, if a configuration ends in 3p 4 , then it needs two electrons to reach a stable state (it takes more energy to lose six, including s-level electrons, so four is easier to lose). And if the configuration ends in 4d 3 , then it needs to lose three electrons to reach a stable state. In addition, half-filled sublevels (s1, p3, d5..) are more stable than, for example, p4 or p2; however, s2 and p6 will be even more stable.
   • When you're dealing with an ion, that means the number of protons is not the same as the number of electrons. The charge of the atom in this case will be shown at the top right (usually) of the chemical symbol. Therefore, an antimony atom with a charge of +2 has the electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 1 . Note that 5p 3 has changed to 5p 1 . Be careful when the configuration of a neutral atom ends at sublevels other than s and p. When you take electrons, you can only take them from valence orbitals (s and p orbitals). Therefore, if the configuration ends with 4s 2 3d 7 and the atom gets +2 charge, then the configuration will end with 4s 0 3d 7 . Please note that 3d 7 not changes, instead electrons of the s-orbital are lost.
   • There are conditions when an electron is forced to "move to a higher energy level." When a sublevel lacks one electron to be half or full, take one electron from the nearest s or p sublevel and move it to the sublevel that needs an electron.
   • There are two options for writing an electronic configuration. They can be written in ascending order of the numbers of energy levels or in the order in which the electron orbitals are filled, as was shown above for erbium.
   • You can also write the electronic configuration of an element by writing only the valence configuration, which is the last s and p sublevel. Thus, the valence configuration of antimony will be 5s 2 5p 3 .
   • Ions are not the same. It's much more difficult with them. Skip two levels and follow the same pattern depending on where you started and how high the number of electrons is.

The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons that have opposite (antiparallel) spins (translated from English as “spindle”), that is, they have properties that can be conditionally represented itself as the rotation of an electron around its imaginary axis: clockwise or counterclockwise. This principle is called the Pauli principle.

If there is one electron in the orbital, then it is called unpaired, if there are two, then these are paired electrons, that is, electrons with opposite spins.

Figure 5 shows a diagram of the division of energy levels into sublevels.

The S-orbital, as you already know, is spherical. The electron of the hydrogen atom (s = 1) is located in this orbital and is unpaired. Therefore, its electronic formula or electronic configuration will be written as follows: 1s 1. In electronic formulas, the energy level number is indicated by the number in front of the letter (1 ...), the sublevel (orbital type) is indicated by the Latin letter, and the number that is written to the upper right of the letter (as an exponent) shows the number of electrons in the sublevel.

For a helium atom, He, having two paired electrons in the same s-orbital, this formula is: 1s 2 .

The electron shell of the helium atom is complete and very stable. Helium is a noble gas.

The second energy level (n = 2) has four orbitals: one s and three p. Second-level s-orbital electrons (2s-orbitals) have a higher energy, since they are at a greater distance from the nucleus than 1s-orbital electrons (n ​​= 2).

In general, for every value of n, there is one s-orbital, but with a corresponding amount of electron energy in it and, therefore, with a corresponding diameter, growing as the value of n increases.

The R-orbital is shaped like a dumbbell or a figure eight. All three p-orbitals are located in the atom mutually perpendicularly along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized again that each energy level (electronic layer), starting from n = 2, has three p-orbitals. As the value of n increases, the electrons occupy p-orbitals located at large distances from the nucleus and directed along the x, y, and z axes.

For elements of the second period (n = 2), first one β-orbital is filled, and then three p-orbitals. Electronic formula 1l: 1s 2 2s 1. The electron is weaker bound to the nucleus of the atom, so the lithium atom can easily give it away (as you obviously remember, this process is called oxidation), turning into a Li + ion.

In the beryllium atom Be 0, 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 - Be 0 is oxidized to the Be 2+ cation.

At the boron atom, the fifth electron occupies a 2p orbital: 1s 2 2s 2 2p 1. Further, the atoms C, N, O, E are filled with 2p orbitals, which ends with the noble gas neon: 1s 2 2s 2 2p 6.

For the elements of the third period, the Sv- and Sp-orbitals are filled, respectively. Five d-orbitals of the third level remain free:

Sometimes, in diagrams depicting the distribution of electrons in atoms, only the number of electrons at each energy level is indicated, that is, they write down the abbreviated electronic formulas of atoms of chemical elements, in contrast to the full electronic formulas given above.

For elements of large periods (fourth and fifth), the first two electrons occupy the 4th and 5th orbitals, respectively: 19 K 2, 8, 8, 1; 38 Sr 2, 8, 18, 8, 2. Starting from the third element of each large period, the next ten electrons will go to the previous 3d and 4d orbitals, respectively (for elements of secondary subgroups): 23 V 2, 8, 11, 2; 26 Tr 2, 8, 14, 2; 40 Zr 2, 8, 18, 10, 2; 43 Tr 2, 8, 18, 13, 2. As a rule, when the previous d-sublevel is filled, the outer (4p- and 5p, respectively) p-sublevel will begin to fill.

For elements of large periods - the sixth and the incomplete seventh - electronic levels and sublevels are filled with electrons, as a rule, as follows: the first two electrons will go to the outer β-sublevel: 56 Ba 2, 8, 18, 18, 8, 2; 87Gr 2, 8, 18, 32, 18, 8, 1; the next one electron (for Na and Ac) to the previous (p-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 energy level from the outside in the 4f and 5f orbitals, respectively, for lanthanides and actinides.

Then the second outside energy level (d-sublevel) will begin to build up again: for elements of secondary subgroups: 73 Ta 2, 8.18, 32.11, 2; 104 Rf 2, 8.18, 32, 32.10, 2 - and, finally, only after the complete filling of the current level with ten electrons will the outer p-sublevel be filled again:

86 Rn 2, 8, 18, 32, 18, 8.

Very often, the structure of the electron shells of atoms is depicted using energy or quantum cells - they write down the so-called graphic electronic formulas. For this record, the following notation is used: each quantum cell is denoted by a cell that corresponds to one orbital; each electron is indicated by an arrow corresponding to the direction of the spin. When writing a graphical electronic formula, two rules should be remembered: the Pauli principle, according to which there can be no more than two electrons in a cell (orbitals, but with antiparallel spins), and F. Hund's rule, according to which electrons occupy free cells (orbitals), are located in they are first one at a time and at the same time have the same spin value, and only then they pair, but the spins in this case, according to the Pauli principle, will already be oppositely directed.

In conclusion, let us once again consider the mapping of the electronic configurations of the atoms of the elements over the periods of the D. I. Mendeleev system. Schemes of the electronic structure of atoms show the distribution of electrons over electronic layers (energy levels).

In a helium atom, the first electron layer is completed - it has 2 electrons.

Hydrogen and helium are s-elements; these atoms have an s-orbital filled with electrons.

Elements of the second period

For all elements of the second period, the first electron layer is filled and the electrons fill the e- and p-orbitals of the second electron layer in accordance with the principle of least energy (first s-, and then p) and the rules of Pauli and Hund (Table 2).

In the neon atom, the second electron layer is completed - it has 8 electrons.

Table 2 The structure of the electron shells of atoms of elements of the second period

The end of the table. 2

Li, Be are β-elements.

B, C, N, O, F, Ne are p-elements; these atoms have p-orbitals filled with electrons.

Elements of the third period

For atoms of elements of the third period, the first and second electron layers are completed; therefore, the third electron layer is filled, in which electrons can occupy the 3s, 3p, and 3d sublevels (Table 3).

Table 3 The structure of the electron shells of atoms of elements of the third period

A 3s-electron orbital is completed at the magnesium atom. Na and Mg are s-elements.

There are 8 electrons in the outer layer (the third electron layer) in the argon atom. As an outer layer, it is complete, 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 p-elements. s- and p-elements form the main subgroups in the Periodic system.

A fourth electron layer appears at the potassium and calcium atoms, and the 4s sublevel is filled (Table 4), since it has a lower energy than the 3d sublevel. To simplify the graphical electronic formulas of the atoms of the elements of the fourth period: 1) we denote the conditionally graphical electronic formula of argon as follows:
Ar;

2) we will not depict the sublevels that are not filled for these atoms.

Table 4 The structure of the electron shells of atoms of the elements of the fourth period

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 the secondary subgroups, they have a pre-external electron layer filled, they are referred to as transition elements.

Pay attention to the structure of the electron shells of chromium and copper atoms. In them, a "failure" of one electron from the 4n- to the 3d sublevel occurs, which is explained by the greater energy stability of the resulting electronic configurations 3d 5 and 3d 10:

In the zinc atom, the third electron layer is complete - all the 3s, 3p and 3d sublevels are filled in it, in total there are 18 electrons on them.

In the elements following zinc, the fourth electron layer, the 4p sublevel, continues to be filled: Elements from Ga to Kr are p-elements.

The outer layer (fourth) of the krypton atom is complete and has 8 electrons. But just in the fourth electron layer, as you know, there can be 32 electrons; the 4d and 4f sublevels of the krypton atom still remain unfilled.

The elements of the fifth period are filling the sublevels in the following order: 5s-> 4d -> 5p. And there are also exceptions associated with the "failure" of electrons, in 41 Nb, 42 MO, etc.

In the sixth and seventh periods, elements appear, that is, elements in which the 4f and 5f sublevels of the third outside electronic layer are being filled, respectively.

The 4f elements are called lanthanides.

5f-elements are called actinides.

The order of filling of electronic sublevels in the atoms of elements of the sixth period: 55 Сs and 56 Ва - 6s-elements;

57 La... 6s 2 5d 1 - 5d element; 58 Ce - 71 Lu - 4f elements; 72 Hf - 80 Hg - 5d elements; 81 Tl - 86 Rn - 6p elements. But even here 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, that is, 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 electronic families or blocks (Fig. 7).

1) s-Elements; the β-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;

2) 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 III-VIII groups;

3) d-elements; the d-sublevel of the preexternal level of the atom is filled with electrons; d-elements include elements of secondary subgroups of groups I-VIII, that is, elements of intercalated decades of large periods located between s- and p-elements. They are also called transition elements;

4) f-elements, the f-sublevel of the third outside level of the atom is filled with electrons; these include lanthanides and actinides.

1. What would happen if the Pauli principle was not respected?

2. What would happen if Hund's rule was not respected?

3. Make diagrams of the electronic structure, electronic formulas and graphic electronic formulas of atoms of the following chemical elements: Ca, Fe, Zr, Sn, Nb, Hf, Ra.

4. Write the electronic formula for element #110 using the symbol for the corresponding noble gas.

5. What is the “failure” of an electron? Give examples of elements in which this phenomenon is observed, write down their electronic formulas.

6. How is the belonging of a chemical element to one or another electronic family determined?

7. Compare the electronic and graphic electronic formulas of the sulfur atom. What additional information does the last formula contain?

Algorithm for compiling 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. By 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 has a maximum of 2 electrons. 1s2, 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.
• Filled in first s- sublevel, then p-, d-b f- sublevels.
• Electrons fill orbitals in ascending order of orbital energy (Klechkovsky's rule).
• Within the sublevel, electrons first occupy free orbitals one at a time, and only after that they form pairs (Hund's rule).
• There cannot be more than two electrons in one orbital (Pauli principle).

Examples.

1. Compose the electronic formula of nitrogen. Nitrogen is number 7 on the periodic table.

2. Compose the electronic formula of argon. In the periodic table, argon is at number 18.

1s 2 2s 2 2p 6 3s 2 3p 6.

3. Compose the electronic formula of chromium. In the periodic table, chromium is number 24.

1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5

Energy diagram of zinc.

4. Compose the electronic formula of zinc. In the periodic table, zinc is number 30.

1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10

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.