What is an electronic structure diagram. Electronic configuration of the atom

Instructions

Electrons in an atom occupy free orbitals in a sequence called a scale: 1s / 2s, 2p / 3s, 3p / 4s, 3d, 4p / 5s, 4d, 5p / 6s, 4d, 5d, 6p / 7s, 5f, 6d, 7p. The orbital can contain two electrons with opposite spins - the directions of rotation.

The structure of electronic shells is expressed using graphical electronic formulas. Use a matrix to write the formula. One cell can contain one or two electrons with opposite spins. Electrons are represented by arrows. The matrix clearly shows that two electrons can be located on the s-orbital, 6 on the p-orbital, 10 on d, and 14 on f.

Write down the sequence number and symbol of the element next to the matrix. In accordance with the energy scale, fill in successive 1s, 2s, 2p, 3s, 3p, 4s levels by inscribing two electrons per cell. It turns out 2 + 2 + 6 + 2 + 6 + 2 \u003d 20 electrons. These levels are completely filled.

You still have five electrons left and an unfilled 3d level. Place the electrons in the cells of the d-sublevel, starting from the left. Place electrons with the same spins in the cells one at a time. If all the cells are filled, starting from the left, add the second electron with the opposite spin. Manganese has five d-electrons, one in each cell.

Electronic graphical formulas clearly show the number of unpaired electrons that determine valence.

note

Remember that chemistry is the science of exceptions. The atoms of the side subgroups of the Periodic Table have an electron "slip". For example, in chromium with serial number 24, one of the electrons from the 4s-level passes into a cell of the d-level. Molybdenum, niobium, etc. have a similar effect. In addition, there is the concept of an excited state of an atom, when paired electrons are unpaired and transfer to neighboring orbitals. Therefore, when drawing up electronic-graphic formulas for the elements of the fifth and subsequent periods of the side subgroup, consult the reference book.

Sources:

• how to compose the electronic formula of a chemical element

Electrons are part of atoms. And complex substances, in turn, are composed of these atoms (atoms form elements) and electrons are divided among themselves. The oxidation state shows which atom took how many electrons for itself, and which gave how many. This indicator can be determined.

You will need

• School chemistry textbooks for grade 8-9 by any author, periodic table, electronegativity table of elements (printed in school chemistry textbooks).

Instructions

To begin with, it is necessary to indicate that the degree is a concept that takes connections for, that is, they do not go deep into the structure. If the element is in a free state, then this is the simplest case - a simple substance is formed, which means that its oxidation state is zero. For example, hydrogen, oxygen, nitrogen, fluorine, etc.

In complex substances, everything is different: electrons are unevenly distributed between atoms, and it is the oxidation state that helps determine the amount of electrons given or received. The oxidation state can be positive or negative. With a plus, electrons are given away, with a minus they are accepted. Some elements retain their oxidation state in various compounds, but many do not differ in this feature. It is necessary to remember an important rule - the sum of the oxidation states is always zero. The simplest example, CO gas: knowing that the oxidation state of oxygen in the overwhelming majority of cases is -2 and using the above rule, you can calculate the oxidation state for C. In sum with -2, zero gives only +2, which means the oxidation state of carbon is +2. Let's complicate the task and take CO2 gas for calculations: the oxidation state of oxygen is still -2, but there are two molecules in this case. Therefore, (-2) * 2 \u003d (-4). The number, in the sum of -4 gives zero, is +4, that is, in this gas it has an oxidation state of +4. An example is more complicated: Н2SO4 - hydrogen has an oxidation state of +1, oxygen has -2. In the given compound there are 2 hydrogen molecules and 4 oxygen molecules, i.e. the charges will be +2 and -8, respectively. In order to get a total of zero, you need to add 6 pluses. This means that the oxidation state of sulfur is +6.

When in a compound it is difficult to determine where is plus, where is minus, an electronegativity table is needed (it is easy to find it in a textbook on general chemistry). Metals often have a positive oxidation state and non-metals a negative one. But for example PI3 - both elements are non-metals. The table indicates that the electronegativity of iodine is 2.6, and that of phosphorus is 2.2. When comparing, it turns out that 2.6 is more than 2.2, that is, electrons are pulled towards iodine (iodine has a negative oxidation state). Following these simple examples, you can easily determine the oxidation state of any element in the compounds.

note

There is no need to confuse metals and non-metals, then the oxidation state will be easier to find and not get confused.

An atom of a chemical element consists of a nucleus and an electron shell. The nucleus is the central part of the atom, in which almost all of its mass is concentrated. Unlike the electron shell, the nucleus has a positive charge.

You will need

• Atomic number of a chemical element, Moseley's law

Instructions

Thus, the charge of the nucleus is equal to the number of protons. In turn, the number of protons in the nucleus is equal to the atomic number. For example, the atomic number of hydrogen is 1, that is, the hydrogen nucleus consists of one proton and has a charge of +1. The atomic number of sodium is 11, the charge of its nucleus is +11.

With alpha decay of a nucleus, its atomic number decreases by two due to the emission of an alpha particle (atomic nucleus). Thus, the number of protons in a nucleus that has undergone alpha decay is also reduced by two.
Beta decay can occur in three different ways. In the case of beta-minus decay, a neutron turns into a proton when an electron and an antineutrino are emitted. Then the charge of the nucleus increases by one.
In the case of beta-plus decay, the proton turns into a neutron, positron and neutrino, the charge of the nucleus decreases by one.
In the case of electron capture, the nuclear charge also decreases by one.

The nuclear charge can also be determined from the frequency of the spectral lines of the characteristic radiation of the atom. According to Moseley's law: sqrt (v / R) \u003d (Z-S) / n, where v is the spectral frequency of the characteristic radiation, R is Rydberg's constant, S is the screening constant, n is the principal quantum number.
So Z \u003d n * sqrt (v / r) + s.

Related Videos

Sources:

• how the charge of the nucleus changes

When creating theoretical and practical work in mathematics, physics, chemistry, a student or schoolchild is faced with the need to insert special symbols and complex formulas. With the Word application from the Microsoft office suite, you can type an electronic formula of any complexity.

Instructions

Click on the Insert tab. On the right, find π, and next to the inscription "Formula". Click on the arrow. A window will appear where you can select a built-in formula like formula quadratic equation.

Click on the arrow and a variety of symbols will appear on the top panel that you may need when writing this particular formula. Once you change it to your liking, you can save it. From now on, it will drop out in the list of built-in formulas.

If you need to transfer the formula to, which later need to be placed on the site, then right-click on the active field with it and select not the professional, but the linear method. In particular, all the same quadratic equation in this case will take the form: x \u003d (- b ± √ (b ^ 2-4ac)) / 2a.

Another option for writing an electronic formula in Word is through the designer. Hold down the Alt and \u003d keys at the same time. You will immediately have a field for writing a formula, and a designer will open in the top panel. Here you can select all the signs that you may need to write an equation and solve any problem.

Some linear notation symbols may be incomprehensible to a reader unfamiliar with computer symbology. In this case, it makes sense to save the most complex formulas or equations in a graphical form. To do this, open the simplest graphics editor Paint: "Start" - "Programs" - "Paint". Then, zoom in on the formula document so that it fills the entire screen. This is necessary for the saved image to have the highest resolution. Press PrtScr on your keyboard, go to Paint and press Ctrl + V.

The electronic structure of the atom can be shown by an electronic formula and an electronic-graphic diagram. In electronic formulas, the energy levels and sublevels are sequentially written in the order of their filling and the total number of electrons on the sublevel. In this case, the state of an individual electron, in particular its magnetic and spin quantum numbers, is not reflected in the electronic formula. In electronic graphic circuits, each electron is "visible" in full, i.e. it can be characterized by all four quantum numbers. Electronic graphing circuits are usually given for external electrons.

Example 1. Write the electronic formula of fluorine, express the state of external electrons with an electronic graphic diagram. How many unpaired electrons are there in an atom of this element?

Decision. The atomic number of fluorine is nine, so there are nine electrons in its atom. In accordance with the principle of least energy, using fig. 7 and taking into account the consequences of the Pauli principle, we write down the electronic formula of fluorine: 1s 2 2s 2 2p 5. For external electrons (the second energy level), we draw up an electronic-graphic diagram (Fig. 8), from which it follows that there is one unpaired electron in the fluorine atom.

Figure: 8. Electronic-graphic diagram of valence electrons of a fluorine atom

Example 2.Draw electronic diagrams of possible states of the nitrogen atom. Which of them reflect the normal state, and which - excited?

Decision.The electronic formula of nitrogen is 1s 2 s 2 2p 3, the formula for the outer electrons is 2s 2 2p 3. The sublevel 2p is incomplete because the number of electrons on it is less than six. Possible variants of the distribution of three electrons on the 2p-sublevel are shown in Fig. nine.

Figure: 9. Electronic-graphic diagrams of possible states of the 2p-sublevel in the nitrogen atom.

The maximum (in absolute value) spin (3/2) corresponds to states 1 and 2, therefore, they are ground, and the rest are excited.

Example 3.Determine the quantum numbers that determine the state of the last electron in the vanadium atom?

Decision. The atomic number of vanadium is Z \u003d 23, therefore, the full electronic formula of the element: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3. The electronic-graphic scheme of external electrons (4s 2 3d 3) is as follows (Fig. 10):

Figure: 10. Electronic-graphic diagram of valence electrons of the vanadium atom

Principal quantum number of the last electron n \u003d 3 (third energy level), orbital l \u003d 2 (sublevel d). The magnetic quantum number for each of the three d-electrons is different: for the first it is –2, for the second –1, for the third - 0. The spin quantum number for all three electrons is the same: m s \u003d + 1/2. Thus, the state of the last electron in the vanadium atom is characterized by quantum numbers: n \u003d 3; l\u003d 2; m \u003d 0; m s \u003d + 1/2.

7. Paired and unpaired electrons

The electrons filling the orbitals in pairs are called paired, and single electrons are called unpaired... Unpaired electrons provide a chemical bond between an atom and other atoms. The presence of unpaired electrons is established experimentally by studying the magnetic properties. Substances with unpaired electrons paramagnetic(are drawn into a magnetic field due to the interaction of the spins of electrons, as elementary magnets, with an external magnetic field). Substances with only paired electrons diamagnetic(the external magnetic field does not act on them). Unpaired electrons are located only on the outer energy level of the atom and their number can be determined by its electronic-graphic scheme.

Example 4.Determine the number of unpaired electrons in a sulfur atom.

Decision. The atomic number of sulfur is Z \u003d 16, therefore, the full electronic formula of the element is: 1s 2 2s 2 2p 6 3s 2 3p 4. The electronic-graphical scheme of external electrons is as follows (Fig. 11).

Figure: 11. Electronic-graphic diagram of valence electrons of the sulfur atom

From the electronic diagram it follows that the sulfur atom has two unpaired electrons.

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 indicates 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 table 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 as a diagram of the distribution of electrons in energy cells.

For iron atoms, such a scheme is as follows:

This diagram clearly shows the fulfillment of the Gund rule. On 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 amendedYES. Mendeleev : the properties of simple bodies, as well as the shapes and properties of compounds of elements, are periodically dependent on the value of the atomic weights of 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 (and not the atomic mass) turned out to be a more accurate argument on which the properties of elements and their compounds depend.

Valence- it is the number of chemical bonds by which one atom is bonded 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 and determines mainly the properties of their atoms. Therefore, these levels are called valence levels. Electrons of these levels, and sometimes of pre-outer 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 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 a given atom interacts. But not all elements freely interact, but practically all elements with oxygen, therefore, the stoichiometric valence can be defined as the doubled number of attached oxygen atoms.

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

When determining the stoichiometric valence of an element according to 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 the other element.

Oxidation statealso characterizes the composition of a substance and is equal to the stoichiometric valence with a plus sign (for a metal or more electropositive element in a molecule) or minus.

1. In simple substances, the oxidation state of the elements is zero.

2. The oxidation state of fluorine in all compounds is -1. The rest of the 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 the compounds has an oxidation state of -2; the exception is 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 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 system, and the lowest is the difference: the 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 cancel each other out 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.

Oxidation degree (oxidative number,) — auxiliary conditional value for recording the processes of oxidation, reduction and redox reactions.

Concept oxidation state is 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 make the bond are completely biased toward more electronegative atoms (that is, assuming that the compound is composed of only ions).

The oxidation state corresponds to the number of electrons that must be attached to a positive ion to reduce it to a neutral atom, or subtract 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, which depend on the structure of the electron shell of the atom, vary by periods and groups of the periodic table. Since in a series of analogous elements the electronic structures are only similar, but not identical, then when passing from one element in a group to another, they observe not a simple repetition of properties, but their more or less clearly expressed regular change.

The chemical nature of an element is due to the ability of its atom to lose or gain electrons. This ability is quantified by the values \u200b\u200bof ionization energies and electron affinity.

Ionization energy (E and) is the minimum amount of energy required for detachment and complete removal of an electron from an atom in the gas phase at T \u003d 0

K without transferring kinetic energy to the liberated electron with the transformation of the atom into a positively charged ion: E + Ei \u003d E + + e-. The ionization energy is a positive value and has the lowest values \u200b\u200bfor alkali metal atoms and the highest for noble (inert) gas atoms.

Electron affinity (Ee) is the energy released or absorbed when an electron attaches to an atom in the gas phase at T \u003d 0

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

E + e- \u003d E- + Ee.

Halogens, especially fluorine (Ee \u003d -328 kJ / mol), have the highest electron affinity.

The values \u200b\u200bof E and Ee are expressed in kilojoules per mole (kJ / mol) or in electron-volts per atom (eV).

The ability of a bound atom to shift the electrons of chemical bonds to 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 the half-sum of the ionization energies and the electron affinity of free atoms ÷ \u003d (Ee + Ei) / 2

In periods, there is a general tendency towards an increase in the ionization energy and electronegativity with an increase in the charge of the atomic nucleus; in groups, these values \u200b\u200bdecrease 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 it enters into, the number and type of neighboring atoms.

Atomic and ionic radii. The sizes of atoms and ions are determined by the size of the electron shell. According to quantum mechanical concepts, the electron shell has no strictly defined boundaries. Therefore, the radius of a free atom or ion can be taken as theoretically calculated distance from the core 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 values \u200b\u200bof the radii of atoms and ions in compounds are usually used, calculated from experimental data. In this case, a distinction is made between covalent and metallic radii of atoms.

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 largest decrease is characteristic of elements of small periods, since their external electronic level is filled. At 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-outer layer. In subgroups, the radii of atoms and ions of the same type generally increase.

The periodic table of elements is a clear example of the manifestation of various kinds of periodicity in the properties of elements, which is observed horizontally (in the 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, while the reducing and metallic properties decrease. So, of all elements of the 3rd period, sodium will be the most active metal and the strongest reducing agent, and chlorine will be the strongest oxidizing agent.

Chemical bond- it is the mutual connection of atoms in a molecule, or crystal lattice, as a result of the action between the atoms of electric forces of attraction.

This is the interaction of all electrons and all nuclei, leading to the formation of a stable, polyatomic system (radical, molecular ion, molecule, crystal).

The chemical bond 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 the molecules there is hydrogen bond, and happen van der Waals interactions.

The main characteristics of the chemical bond include:

- bond length - this 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;

- 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 the 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 its breaking, kJ / mol.

Covalent bond - Chemical bond, formed by socialization of a pair of electrons by 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 instrument 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 MFM:

1. The 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 bond is stronger, 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 socializing the valence electrons of two neutral atoms. Each atom gives one unpaired electron to a common electron pair:

Figure: 7. Exchange mechanism of covalent bond formation: and - non-polar; b - polar

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

Connections, educatedby donor-acceptor mechanism, refer to complex compounds

Figure: 8. Donor-acceptor mechanism of covalent bond formation

The 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.

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 planar 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 is carried out have a tetrahedral geometry (CH 4 , NH 3).

Figure: 10. Types of hybridizations of valence orbitals: a - sp-hybridization of valence orbitals; b - sp 2 -hybridization of valence orbitals; in - sp 3-hybridization of valence orbitals

2. The structure of nuclei and electron shells of atoms

2.7. Distribution of electrons in an atom

The state of electrons in an atom is indicated using a specific notation. For example, for a helium atom we have:

The distribution of electrons in an atom is indicated by:

and) electronic circuits, in which only the number of electrons in each layer is noted. For example: Mg 2e, 8e, 2e; Cl 2e, 8e, 7e.

Often used graphic electronic circuits, for example, for a chlorine atom:

b) electronic configurations; in this case, the number of the layer (level), the nature of the sublevels, and the number of electrons on them are shown. For instance:
Mg 1s 2 2s 2 2p 6 3s 2;

at) electronic-graphic schemes, on which the orbitals are depicted, for example, in the form of a cell, and the electrons - by arrows (Fig. 2.6).

Figure: 2.6. Electronic diagram for the magnesium atom

In addition to the complete formulas for electronic configurations, abbreviated ones are widely used. In this case, the noble gas portion of the electronic configuration is indicated by the noble gas symbol in square brackets. For example: 12 Mg3s 2, 19 K4s 1.

There are certain principles and rules for filling energy levels and sublevels with electrons:

1. The principle of the minimum total energy of the atom, according to which the population of the AO with electrons occurs so that the total energy of the atom is minimal. The following sequence of filling the AO has been established experimentally:

1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p ....

2. One AO \u200b\u200bcan contain no more than two electrons, and their spins in this case must be antiparallel.

3. Within a given energy sublevel, electrons fill the AO gradually, first one by one (first, all vacant, and then two), and the orientation of all unpaired electrons should be the same, i.e. such

but not like that

In almost any atom, only s - and p -AO are external (Fig. 2.7), therefore there can be no more than eight electrons on the outer electron layer of any atom... The outer electron layer containing eight electrons (in the case of helium, two) is called complete.

Figure: 2.7. Electronic-graphic schemes for atoms K (a) and S (b)

The values \u200b\u200bof the energies of different energy sublevels for different atoms are not constant, but depend on the charge of the nucleus Z of the atom of the element: for atoms of elements with Z \u003d 1–20 Е 3 d\u003e E 4 s and Е 3 d\u003e E 4 p; vice versa for atoms of elements with Z ≥ 21: E 3 d< E 4 s и Е 3 d < E 4 p (рис. 2.8). Кроме того, чем больше Z , тем меньше различаются подуровни по энергии, а кривые, выражающие зависимость энергии подуровней от Z , пересекаются.

Figure: 2.8. Diagram of energy sublevels of atoms of elements with Z \u003d 1–20 (a), Z ≥ 21 (b)

The electronic configurations of the atoms (ground state) of K and Ca are as follows (see Fig. 2.8):

19 K: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1,

20 Ca: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2.

Starting with scandium (Z \u003d 21), the 3d sublevel is filled, and 4s electrons remain in the outer layer. The general electronic formula of atoms of elements from Sc to Zn is 3d 1-10 4s 1-2. For instance:

21 Sс: 3d 1 4s 2,

25 Mn: 3d 5 4s 2,

28 Ni: 3d 8 4s 2.

30 Zn: 3d 10 4s 2.

For chromium and copper, a slip (dip) of the 4s electron to the 3d sublevel is observed: Cr - 3d 5 4s 1, Cu - 3d 10 4s 1. Such a slip from ns - to the (n - 1) d-sublevel is also observed in atoms of other elements (Mo, Ag, Au, Pt) and is explained by the closeness of the energies of the ns - and (n - 1) d-sublevels, as well as by the stability of half and completely filled d-sublevels.

Further in the 4th period after 10 d -elements follow from Ga (3d 10 4s 2 4p 1) to Kr (3d 10 4s 2 4p 6) p -elements.

The formation of cations of d -elements is associated with the loss of first external ns -, then (n - 1) d -electrons, for example:

Ti: 3d 2 4s 2 → - 2 e - Ti 2+: 3d 2 → - 1 e - Ti 3+: 3d 1

Mn: 3d 5 4s 2 → - 2 e - Mn 2+: 3d 5 → - 2 e - Mn 4+: 3d 3

Note that in the formulas of electronic configurations, it is customary to write down first all electrons with a lower value of n, and then proceed to indicate electrons with a higher value of the principal quantum number. Therefore, the filling order and the order of writing the energy sublevels for 3d elements do not coincide. For example, in the electronic formula of the scandium atom, the 3d orbital is indicated up to the 4s orbital, although the 4s orbital is filled earlier.

A natural question arises: why is the 4s-sublevel filled earlier in atoms of 3d -elements, although its energy is higher than the energy of the 3d -sublevel? Why, for example, does the Sc atom not have the electronic configuration 1s 2 2s 2 2p 6 3s 2 3p 6 3d 3 in the ground state?

This happens because the ratio of the energies of various electronic states of an atom does not always coincide with the ratio of the energies of individual energy sublevels. The energy of the 4s-sublevel for 3d -elements is greater than the energy of the 3d -sublevel, but the energy of the state
3d 1 4s 2 is less than the energy of the 3d 3 state.

This is explained by the fact that the electron-electron repulsion, and, accordingly, the energy of the entire state for the configuration ... 3d 3 (with three electrons at the same energy sublevel) is greater than for the configuration ... 3d 1 4s 2 (with three electrons, located at different energy levels).

DEFINITION

Electronic formula (configuration) of an atom of a chemical element shows the arrangement of electrons on the electron shells (levels and sublevels) in an atom or molecule.

Most often, electronic formulas are written for atoms in the ground or excited state and for ions.

There are several rules that must be taken into account when drawing up the electronic formula of an atom of a chemical element. This is Pauli's principle, Klechkovsky's rules or Hund's rule.

Drawing up an electronic and electronic-graphic formula

When drawing up an electronic formula, it should be borne in mind that the number of the period of a chemical element determines the number of energy levels (shells) in the atom, and its ordinal number determines the number of electrons.

According to the Klechkovsky rule, the filling of energy levels occurs in the increasing order of the sum of the principal and orbital quantum numbers (n + l), and for equal values \u200b\u200bof this sum, in the increasing order of n:

1s< 2s < 2p < 3s < 3p < 4s ≈ 3d < 4p < 5s ≈ 4d < 5p < 6s ≈ 5d ≈ 4f < 6p и т.д.

Thus, the value n + l \u003d 5 corresponds to the energy sublevels 3d (n \u003d 3, l \u003d 2), 4d (n \u003d 4, l \u003d 1), and 5s (n \u003d 5, l \u003d 0). The first of these sublevels is filled with the one with the lower value of the main quantum number.

The behavior of electrons in atoms obeys the principle of exclusion formulated by the Swiss scientist W. Pauli: there cannot be two electrons in an atom, which would have all four quantum numbers the same. According to pauli principle, in one orbital, characterized by certain values \u200b\u200bof the three quantum numbers (principal, orbital and magnetic), there can be only two electrons that differ in the value of the spin quantum number. Pauli's principle implies consequence: the maximum possible number of electrons at each energy level is equal to twice the square of the principal quantum number.

The electronic formula of the atom is depicted as follows: each energy level corresponds to a certain principal quantum number n, denoted by an Arabic numeral; each number is followed by a letter corresponding to the energy sublevel and denoting the orbital quantum number. The superscript next to the letter indicates the number of electrons in the sublevel. For example, the electronic formula for the sodium atom is:

11 N 1s 2 2s 2 2p 6 3s 1.

When filling the energy sublevels with electrons, it is also necessary to observe hund's rule: in this sublevel, electrons tend to occupy energy states in such a way that the total spin is maximum, which is most clearly reflected in the compilation of electronic-graphic formulas.

Electronic graphic formulas usually depicted for valence electrons. In such a formula, all electrons are marked with arrows, and cells (squares) - orbitals. One cell cannot contain more than two electrons. Let's consider vanadium as an example. First, we write down the electronic formula and determine the valence electrons:

74 W) 2) 8) 18) 32) 12) 2;

1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4f 14 5s 2 5p 6 5d 4 6s 2 .

The outer energy level of a tungsten atom contains 6 electrons, which are valence. The energy diagram of the ground state takes the following form:

Examples of problem solving

EXAMPLE 1