Cl electron distribution. Distribution of electrons in atoms

If identical particles have the same quantum numbers, then their wave function is symmetric with respect to the permutation of particles. It follows that two identical fermions included in the same system cannot be in the same states, because for fermions the wave function must be antisymmetric. Summarizing the experimental data, W. Pauli formed principle exceptions , according to which fermion systems occur in nature only in states,described by antisymmetric wave functions(quantum mechanical formulation of the Pauli principle).

From this position follows a simpler formulation of the Pauli principle, which he introduced into quantum theory (1925) even before the construction of quantum mechanics: in a system of identical fermions any two of them cannot simultaneously be in the same state . Note that the number of identical bosons in the same state is not limited.

Let us recall that the state of an electron in an atom is uniquely determined by the set four quantum numbers :

· main n ;

· orbital l , usually these states are designated 1 s, 2d, 3f;

magnetic();

· magnetic spin ().

The distribution of electrons in an atom occurs according to the Pauli principle, which can be formulated for an atom in its simplest form: the same atom cannot have more than one electron with the same set of four quantum numbers: n, l, , :

Z (n, l, , ) = 0 or 1,

Where Z (n, l, , ) - the number of electrons in a quantum state, described by a set of four quantum numbers: n, l. . . Thus, the Pauli principle states that two electrons ,bound in the same atom differ in meaning ,at least ,one quantum number .

The maximum number of electrons in states described by a set of three quantum numbers n, l And m, and differing only in the orientation of the electron spins is equal to:

,

(8.2.1)

because the spin quantum number can take only two values ​​1/2 and –1/2.

The maximum number of electrons in states defined by two quantum numbers n And l:

.

(8.2.2)

In this case, the vector of the orbital angular momentum of the electron can take in space (2 l+ 1) different orientations (Fig. 8.1).

The maximum number of electrons in states determined by the value of the principal quantum number n, equals:

.

(8.2.3)

Collection of electrons in a multi-electron atom,having the same principal quantum number n,called electron shell or layer .

In each shell, electrons are distributed according to subshells , corresponding to this l.

Region of space,in which there is a high probability of detecting an electron, called subshell or orbital . The main types of orbitals are shown in Fig. 8.1.

Since the orbital quantum number takes values ​​from 0 to , the number of subshells is equal to the ordinal number n shells. The number of electrons in a subshell is determined by the magnetic and magnetic spin quantum numbers: the maximum number of electrons in a subshell with a given l equals 2(2 l+ 1). Shell designations, as well as the distribution of electrons across shells and subshells are given in Table. 1.

Table 1

Principal quantum number n

Shell symbol

Maximum number of electrons in shell

Orbital quantum number l

Subshell symbol

Maximum number electrons in subshell

The distribution is characterized by the following rules:

Pauli principle;

Hund's rule;

the principle of least energy and Klechkovsky's rule.

By Pauli principle An atom cannot have two or more electrons with the same value of all four quantum numbers. Based on the Pauli principle, the maximum capacity of each energy level and sublevel can be determined.

	Sublevel, ℓ	Sublevel designation	Magnetic quantum number, m	Spin quantum number,s
			3, -2, -1, 0, 1, 2, 3

Thus, maximum number of electrons per:

s -sublevel – 2,

p - sublevel – 6,

d -sublevel – 10,

f -sublevel – 14.

Within the quantum level n, an electron can take on the values ​​of 2n 2 different states, which was established experimentally using spectral analysis.

Hund's rule : in each sublevel, electrons strive to occupy the maximum number of free energy cells so that the total spin has the greatest value.

For example:

right wrong wrong

3p 3:

s = +1/2+1/2+1/2=1.5 s =-1/2+1/2+1/2=0.5 s = -1/2+1/2-1/2 =-0.5

The principle of least energy and Klechkovsky's rule: electrons primarily occupy quantum orbitals with minimal energy. Since the energy reserve in an atom is determined by the value of the sum of the principal and orbital quantum numbers (n + ℓ), electrons first occupy the orbitals for which the sum (n + ℓ) is the smallest.

For example: the sum (n + ℓ) for the 3d sublevel is n = 3, l = 2, therefore (n + ℓ) = 5; for the 4s sublevel: n = 4, ℓ = 0, therefore (n + ℓ ) = 4. In this case, the 4s sublevel is filled first and only then the 3d sublevel.

If the total energy values ​​are equal, then the level that is closest to the nucleus is populated.

For example: for 3d: n = 3, ℓ = 2 , (n + ℓ) = 5 ;

for 4p: n = 4, ℓ = 1, (n + ℓ) = 5.

Since n = 3 < n = 4, 3d will be populated with electrons earlier than 4 p.

Thus, sequence of filling levels and sublevels with electrons in atoms:

1 s 2 <2 s 2 <2 p 6 <3 s 2 <3 p 6 <4 s 2 <3 d 10 <4 p 6 <5 s 2 <4 d 10 <5 p 6 <6 s 2 <5 d 10 ≈ 4 f 14 <6 p 6 <7s 2 …..

Electronic formulas

An electron formula is a graphical representation of the distribution of electrons across levels and sublevels in an atom. There are two types of formulas:

When writing, only two quantum numbers are used: n and ℓ. The main quantum number is indicated by a number before the letter designation of the sublevel. The orbital quantum number is indicated by the letter s, p, d or f. The number of electrons is indicated by a number as an exponent.

For example: +1 H: 1s 1 ; +4 Be: 1s 2 2s 2 ;

2 He: 1s 2 ; +10 Ne: 1s 2 2s 2 2p 6 ;

3 Li: 1s 2 2s 1 ; +14 Si: 1s 2 2s 2 2p 6 3s 2 3p 6 .

That is, the sequence is observed

1 s 2 <2 s 2 <2 p 6 <3 s 2 <3 p 6 <4 s 2 <3 d 10 <4 p 6 <5 s 2 <4 d 10 <5 p 6 <6 s 2 <5 d 10 ≈ 4 f 14 <6 p 6 <7s 2 …..

graphical electronic formula - all 4 quantum numbers are used - this is the distribution of electrons across quantum cells. The main quantum number is depicted on the left, the orbital number is represented by the letter below, the magnetic number is the number of cells, and the spin number is the direction of the arrows.

For example:

8 O:…2s 2 2p 4

The graphical formula is used to write only valence electrons.

Let's consider compiling electronic formulas of elements by periods.

The first period contains 2 elements in which the first quantum level and the s-sublevel are completely populated with electrons (the maximum number of electrons per sublevel is 2):

2 He: n=1 1s 2

Elements whose s-sublevel is filled last are classified as s -family and call s -elements .

For period II elements, the II quantum level, s- and p-sublevels are being filled (the maximum number of electrons at the p-sublevel is 8).

3 Li: 1s 2 2s 1 ; 4 Be: 1s 2 2s 2 ;

5 B: 1s 2 2s 2 2p 1 ; 10 Ne: 1s 2 2s 2 2p 6

Elements whose p-sublevel is filled last are classified as p-family and call p-elements .

Elements of the III period begin to form the III quantum level. In Na and Mg, the 3s sublevel is populated with electrons. For elements from 13 Al to 18 Ar, the 3p sublevel is populated; The 3d sublevel remains unfilled, since it has a higher energy level than the 4s sublevel and is not filled in elements of the III period.

The 3d sublevel begins to be filled in elements of the IV period, and the 4d - in elements of the V period (in accordance with the sequence):

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 ;

21 Sc: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 1 ; 25 Mn: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 ;

33 As: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 3 ; 43 Tc: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 5

Elements whose d-sublevel is filled last are classified as d -family and call d -elements .

4f is filled in only after element 57 of the VI period:

57 La: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 5s 2 4d 10 5p 6 6s 2 5d 1 ;

58 Ce: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 5 5s 2 4d 10 5p 6 6s 2 5d 1 4f 1 ;

The population of the V quantum level by electrons proceeds similarly to the IV period. Thus, the previously shown sequence of population of levels and sublevels by electrons is observed:

6s 2 5d 10 4f 14 6p 6

the population of a new quantum level by electrons always begins with the s-sublevel. For elements of a given period, only the s and p sublevels of the external quantum level are populated by electrons;

the population of the d-sublevel is delayed by period I; The 3d sublevel is filled in for elements of the IV period, the 4d sublevel is filled in for elements of the V period, etc.;

the population of the f sublevel by electrons is delayed by 2 periods; The 4f sublevel is populated in elements of the VI period, the 5f sublevel is populated in elements of the VII period, etc.

Electrons are distributed among sublevels, forming clouds of a certain shape around the nucleus; this distribution depends on the amount of their energy, that is, the closer the electron is to the nucleus of the atom, the less its amount of energy.

Electrons tend to occupy a position corresponding to the minimum energy value and are located around the nucleus according to the Pauli principle. As is known from previous topics, the largest number of electrons that can be located in each electronic layer is determined by the formula N = 2n 2. The first electron layer or layer K is located at the closest distance from the nucleus of the atom and has n=1. In accordance with this, N = 2-1 2 = 2 electrons move on this layer. The second electron layer can accommodate 8, the third - 18, and the fourth - 32 electrons.

In the outer electronic layers of all elements (except for elements of the 1st period) there are no more than eight electrons. The outer electron layers of noble gases (with the exception of helium) are filled with eight electrons, so these gases are chemically stable.

At the outer energy level of the elements of the main subgroup of the periodic table, the number of electrons is equal to the group number. The number of electrons in the outer layer of elements of the side subgroup does not exceed two; when moving from one element to the second, the attracted electrons move from the outer layer to the inner one, since the outer layer is replenished with ns 2 nр 6 electrons, and the joining electrons occupy the nd sublevel.

Thus, the manganese atom has the following structure: Mn(+25) 2, 8, 13, 2, and its electronic formula: 1s 2 2s 2 2p 6 3s 2 3p 6 3d 5 4s 2.

According to the Pauli principle, any atom cannot have two electrons with the same quantum numbers.

Consequently, in each orbital of an atom, the value of three quantum numbers - n, l, m (principal, orbital and magnetic) can be the same, but the spin quantum numbers (s) differ, that is, there are electrons with opposite spins.

The replenishment of sublevels with electrons was clarified using the rule of V.M. Klechkovsky (1900-1972) according to which electrons fill energy sublevels in the following order:

The order of filling cells (cells) of energy levels with electrons obeys Hund's rule. First, the 2p cells are filled with six electrons. The next electron, according to the Klechkovsky rule, goes into the 3s energy sublevel:

19. Klechkovsky's rule reads:

The n + l rule was proposed in 1936 by the German physicist E. Madelung; in 1951 it was again formulated by V. M. Klechkovsky.

The electron shell of an atom is a region of space where electrons are likely to be located, characterized by the same value of the principal quantum number n and, as a consequence, located at close energy levels. The number of electrons in each electron shell does not exceed a certain maximum value.

The order of filling electron shells (orbitals with the same value of the principal quantum number n) is determined by the Klechkovsky rule, the order of filling orbitals within the same sublevel with electrons (orbitals with the same values ​​of the principal quantum number n and orbital quantum number l) is determined by Hund’s Rule.

20.Atomic nucleus- the central part of the atom, in which the bulk of its mass is concentrated (more than 99.9%). The nucleus is positively charged; the charge of the nucleus is determined by the chemical element to which the atom belongs. The sizes of the nuclei of various atoms are several femtometers, which is more than 10 thousand times smaller than the size of the atom itself.

The atomic nucleus consists of nucleons - positively charged protons and neutral neutrons, which are connected to each other through the strong interaction

The number of protons in a nucleus is called its charge number - this number is equal to the atomic number of the element to which the atom belongs in Mendeleev’s table (Periodic Table of Elements). The number of protons in the nucleus determines the structure of the electron shell of a neutral atom and, thus, the chemical properties of the corresponding element. The number of neutrons in a nucleus is called its isotopic number. Nuclei with the same number of protons and different numbers of neutrons are called isotopes. Nuclei with the same number of neutrons but different numbers of protons are called isotones. The terms isotope and isotone are also used to refer to atoms containing these nuclei, as well as to characterize non-chemical varieties of a single chemical element. The total number of nucleons in a nucleus is called its mass number () and is approximately equal to the average mass of an atom shown in the periodic table. Nuclides with the same mass number but different proton-neutron composition are usually called isobars.

nuclear reaction- the process of transformation of atomic nuclei that occurs during their interaction with elementary particles, gamma rays and with each other. A nuclear reaction is the process of interaction of an atomic nucleus with another nucleus or elementary particle, accompanied by a change in the composition and structure of the nucleus and the release of secondary particles or γ-quanta. The nuclear reaction was first observed by Rutherford in 1919, bombarding the nuclei of nitrogen atoms with α particles; it was detected by the appearance of secondary ionizing particles that had a range in the gas greater than that of the α particles and were identified as protons. Subsequently, photographs of this process were obtained using a cloud chamber.

According to the mechanism of interaction, nuclear reactions are divided into two types:

· reactions with the formation of a compound nucleus are a two-stage process that occurs at a not very high kinetic energy of colliding particles (up to about 10 MeV).

· direct nuclear reactions that take place in the nuclear time required for a particle to cross the nucleus. This mechanism mainly manifests itself at high energies of bombarding particles.

Only a small part of nuclides are stable. In most cases, nuclear forces are unable to ensure their continued integrity, and the nuclei sooner or later disintegrate. This phenomenon is called radioactivity.

Radioactivity

Radioactivity is the ability of an atomic nucleus to spontaneously decay by emitting particles. Radioactive decay is characterized by the lifetime of the radioactive isotope, the type of particles emitted, and their energies.
The main types of radioactive decay are:

• α-decay – emission of an α-particle by an atomic nucleus;
• β-decay – emission of an electron and an antineutrino, a positron and a neutrino by an atomic nucleus, absorption of an atomic electron by the nucleus with the emission of a neutrino;
• γ-decay – emission of γ-quanta by an atomic nucleus;

· spontaneous fission - the disintegration of an atomic nucleus into two fragments of comparable mass.

21. periodic system and periodic law By the beginning of the 19th century. About 30 elements were known, by the middle of the 19th century - about 60. As the number of elements accumulated, the task of systematizing them arose. Such attempts before D.I. Mendeleev was no less than fifty; The systematization was based on: atomic weight (now called atomic mass), chemical equivalent, and valence. Approaching the classification of chemical elements metaphysically, trying to systematize only the elements known at that time, none of D.I. Mendeleev’s predecessors could discover the universal interconnection of elements or create a single harmonious system reflecting the law of development of matter. This important problem for science was brilliantly solved in 1869 by the great Russian scientist D.I. Mendeleev, who discovered the periodic law.
Mendeleev’s systematization was based on: a) atomic weight and b) chemical similarity between elements. The most striking expression of the similarity of the properties of elements is their identical highest valency. Both atomic weight (atomic mass) and the highest valence of an element are quantitative, numerical constants convenient for systematization.
Having arranged all 63 elements known at that time in a row in order of increasing atomic masses, Mendeleev noticed the periodic repeatability of the properties of elements at unequal intervals. As a result, Mendeleev created the first version of the periodic table.
The regular nature of the change in the atomic masses of elements along the verticals and horizontals of the table, as well as the empty spaces formed in it, allowed Mendeleev to boldly predict the presence in nature of a number of elements that were not yet known to science at that time and even outline their atomic masses and basic properties based on the expected position elements in the table. This could only be done on the basis of a system that objectively reflects the law of development of matter. The essence of the periodic law D.I. Mendeleev formulated in 1869: “The properties of simple bodies, as well as the forms and properties of compounds of elements are periodically dependent on the magnitude of the atomic weights (mass) of the elements.”

The design of the modern periodic table, in principle, differs little from the version of 1871. The symbols of the elements in the periodic system are arranged in vertical and horizontal columns. This leads to the unification of elements into groups, subgroups, periods. Each element occupies a specific cell in the table. Vertical graphs are groups (and subgroups), horizontal graphs are periods (and series).

Covalent bond

The bond that arises when electrons interact with the formation of generalized electron pairs is called covalent.

If the interacting atoms have equal electronegativity values, the shared electron pair belongs equally to both atoms, that is, it is located at an equal distance from both atoms. This covalent bond is called non-polar. It occurs in simple non-metallic substances: H22, O22, N22, Cl22, P44, O33.

When atoms that have different electronegativity values, such as hydrogen and chlorine, interact, the shared electron pair is shifted towards the atom with higher electronegativity, that is, towards chlorine.

The chlorine atom acquires a partial negative charge, and the hydrogen atom acquires a partial positive charge. This is an example polar covalent bond.

Properties of covalent bonds

The characteristic properties of a covalent bond - directionality, saturation, polarity, polarizability - determine the chemical and physical properties of organic compounds.

Communication direction determines the molecular structure of organic substances and the geometric shape of their molecules. The angles between two bonds are called bond angles.

Saturability- the ability of atoms to form a limited number of covalent bonds. The number of bonds formed by an atom is limited by the number of its outer atomic orbitals.

The polarity of the bond is due to the uneven distribution of electron density due to differences in the electronegativity of the atoms. On this basis, covalent bonds are divided into non-polar and polar.

The polarizability of a bond is expressed in the displacement of the bond electrons under the influence of an external electric field, including that of another reacting particle. Polarizability is determined by electron mobility. Electrons are more mobile the further they are from the nuclei.

The polarity and polarizability of covalent bonds determines the reactivity of molecules towards polar reagents.

23. Ionic bond- a chemical bond formed between atoms with a large difference in electronegativity, in which the shared electron pair is completely transferred to the atom with a higher electronegativity.
Since an ion can attract ions of the opposite sign to itself in any direction, an ionic bond differs from a covalent bond in its non-directional nature.

The interaction of two ions of opposite sign with each other cannot lead to complete mutual compensation of their force fields. Therefore, they can attract other ions of the opposite sign, that is, the ionic bond is unsaturated.

24. Metal connection- a chemical bond between atoms in a metal crystal, arising due to the sharing of their valence electrons.

Metal connection- communication between positive ions in metal crystals, carried out due to the attraction of electrons moving freely throughout the crystal. According to their position in the periodic table, metal atoms have a small number of valence electrons. These electrons are rather weakly bound to their nuclei and can easily break away from them. As a result, positively charged ions and free electrons appear in the crystal lattice of the metal. Therefore, in the crystal lattice of metals there is great freedom of movement of electrons: some of the atoms will lose their electrons, and the resulting ions can accept these electrons from the “electron gas”. As a consequence, the metal represents a number of positive ions localized in certain positions of the crystal lattice, and a large number of electrons moving relatively freely in the field of positive centers. This is an important difference between metallic bonds and covalent bonds, which have a strict orientation in space.

A metal bond also differs from a covalent bond in strength: its energy is 3-4 times less than the energy of a covalent bond.

Hydrogen bond

A hydrogen atom connected to a fluorine, oxygen or nitrogen atom (less commonly, chlorine, sulfur or other non-metals) can form another additional bond. This discovery, made in the eighties of the nineteenth century, is associated with the names of Russian chemists M.A. Ilyinsky and N.N. Beketova. It was found that some hydrogen-containing groups of atoms often form stable chemical bonds with electronegative atoms that are part of another or the same molecule. This chemical bond is called a hydrogen bond.

A hydrogen bond is an interaction between two electronegative atoms of the same or different molecules through a hydrogen atom: A−H... B (a line indicates a covalent bond, three dots indicate a hydrogen bond).

A hydrogen bond is caused by the electrostatic attraction of a hydrogen atom (carrying a positive charge δ+) to an atom of an electronegative element having a negative charge δ−. In most cases, it is weaker than covalent, but is significantly stronger than the usual attraction of molecules to each other in solid and liquid substances. Unlike intermolecular interactions, a hydrogen bond has the properties of directionality and saturation, therefore it is often considered one of the types of covalent chemical bonds. It can be described using the molecular orbital method as a three-center, two-electron bond.

One of the signs of a hydrogen bond can be the distance between a hydrogen atom and another atom forming it. It must be less than the sum of the radii of these atoms. More common are asymmetric hydrogen bonds, in which the distance H...B is greater than A−B. However, in rare cases (hydrogen fluoride, some carboxylic acids) the hydrogen bond is symmetrical. The angle between atoms in the A−H...B fragment is usually close to 180 o. The strongest hydrogen bonds are formed with the participation of fluorine atoms. In a symmetric ion, the energy of the hydrogen bond is 155 kJ/mol and is comparable to the energy of the covalent bond. The hydrogen bond energy between water molecules is already noticeably lower (25 kJ/mol).

26. Thermal effect of a chemical reaction or a change in the enthalpy of a system due to the occurrence of a chemical reaction - the amount of heat attributed to the change in a chemical variable received by the system in which a chemical reaction took place and the reaction products took on the temperature of the reactants.

For the thermal effect to be a quantity that depends only on the nature of the ongoing chemical reaction, the following conditions must be met:

· The reaction must proceed either at constant volume Q v (isochoric process) or at constant pressure Q p (isobaric process).

· No work is performed in the system, except for the expansion work possible at P = const.

If the reaction is carried out under standard conditions at T = 298.15 K = 25 ˚C and P = 1 atm = 101325 Pa, the thermal effect is called the standard thermal effect of the reaction or the standard enthalpy of the reaction ΔH r O. In thermochemistry, the standard heat of reaction is calculated using standard enthalpies of formation.

Hess's Law (1841)

The thermal effect (enthalpy) of a process depends only on the initial and final states and does not depend on the path of its transition from one state to another.

28. Rate of chemical reaction- change in the amount of one of the reacting substances per unit of time in a unit of reaction space. Is a key concept in chemical kinetics. The rate of a chemical reaction is always a positive value, therefore, if it is determined by the starting substance (the concentration of which decreases during the reaction), then the resulting value is multiplied by −1.

In 1865, N. N. Beketov and in 1867, Guldberg and Waage formulated the law of mass action: the rate of a chemical reaction at each instant of time is proportional to the concentrations of reagents raised to powers equal to their stoichiometric coefficients

For elementary reactions, the exponent of the concentration of each substance is often equal to its stoichiometric coefficient; for complex reactions this rule is not observed. In addition to concentration, the following factors influence the rate of a chemical reaction:

the nature of the reactants,

presence of a catalyst,

temperature (van't Hoff rule, Arrhenius equation),

· pressure,

· surface area of ​​reacting substances.

If we consider the simplest chemical reaction A + B → C, we will notice that the instantaneous rate of a chemical reaction is not constant

29. Law of mass action. In 1865, Professor N.N. Beketov was the first to hypothesize a quantitative relationship between the masses of reactants and the reaction time. This hypothesis was confirmed in the law of mass action, which was established in 1867 by two Norwegian chemists K. Guldberg and P. Waage. The modern formulation of the law of mass action is as follows:

At constant temperature, the rate of a chemical reaction is directly proportional to the product of the concentrations of the reacting substances, taken in powers equal to the stoichiometric coefficients in the reaction equation.

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Guidelines

17. Taking into account the described patterns, consider the state and distribution of electrons across energy levels and orbitals for potassium atoms ( Z= 19) and scandium ( Z = 21).

Solution

1) The element preceding potassium in PSCE is argon ( Z= 18) has the following electron distribution:

a) by atomic levels:

b) according to the orbitals of the atom:

Electronic formula of the argon atom:

Electron graphic formula of the argon atom:

When distributing electrons in the K atom in accordance with the Klechkovsky rule, preference is given to orbital 4 s(sum of quantum numbers n + l equal to: 4 + 0 = 4) compared to orbital 3 d(sum of quantum numbers n + l equal to: 3 + 2 = 5) as the orbital having the minimum value n + l. Therefore, for the potassium atom, the distribution of electrons over orbitals (electron graphic formula) has the form (see paragraph 16 of the guidelines):

Potassium belongs to s-elements with the following electronic formula (configuration) of the atom:

The electron energy level distribution for the K atom is shown below:

2) The element preceding scandium in PSCE is calcium ( Z= 20) has the following electron distribution:

a) by atomic levels:

b) according to the orbitals of the atom:

Electronic formula of the calcium atom:

From orbitals 3 d (n + l equals: 3 + 2 = 5) and 4 p (n + l equals: 4 + 1 = 5) when distributing electrons in a scandium atom among orbitals, preference should be given to 3 d-orbital as having the minimum value n= 3 for the same sums of quantum numbers ( n + l) equal to five. Therefore, scandium belongs to d-elements, and its atom is characterized by the following distribution of electrons among orbitals:

Electronic formula of the scandium atom:

The electron energy level distribution for the Sc atom is depicted below:

18. Complete the drawing to show the appearance of one s-orbitals and three r-orbitals oriented along the axes.

Table 5

Electron distribution
by quantum levels and sublevels

Shell	Energy level n	Energy sublevel l	Magnetic number m	Number orbitals	Limit number electrons
K	1	0(s)	0	1	2
L	2	0(s) 1 (p)	+1, 0, –1	1 3 4	2 6 8
M	3	0(s) 1 (p) 2(d)	0 1, 0, –1 +2, +1, 0, –1, –2	1 3 5 9	2 6 10 18
N	4	0(s) 1 (p) 2(d) 3(f)	0 +1, 0, –1 +2, +1, 0, –1, –2 +3, +2, +1, 0, –1, –2, –3	1 3 5 7 16	2 6 10 14 32

20. For the sequence of filling the energy levels of atoms, see table. 6.

21. The number of elements in a period of D.I. Mendeleev’s table is determined by the formulas:

a) for odd periods:

Ln = (n + 1) 2 /2,

b) for even periods:

Ln = (n + 2) 2 /2,

Where Ln– number of elements in the period, n– period number.

Define the number of elements in each period of D.I. Mendeleev’s PSHE.

Explain:

a) the resulting numerical pattern from the standpoint of the state of electrons in atoms and their distribution among energy levels;

b) division of groups of elements into main and secondary subgroups;

c) the predetermination of the number of main and secondary subgroups in D.I. Mendeleev’s PSHE from the point of view of the theory of atomic structure.

Check in the future, your conclusions on Appendix 1 (P-21).

22. The strict periodicity of the arrangement of elements in D.I. Mendeleev’s PSHE is fully explained by the sequential filling of the energy levels of atoms (see paragraph 20 above). The strengthening of the position of the periodic law based on the patterns of changes in the electronic structure of atoms of elements, first predicted by N. Bohr, was facilitated by the discovery of the 72nd element. Chemists searched for the then-undiscovered element among minerals containing rare earth elements, based on the incorrect premise that 15 elements should be classified as lanthanides.

By analogy with transition elements, the number of lanthanides (elements No. 58–71) should be equal to the difference between the maximum numbers of electrons per N And M energy levels
(32 – 18 = 14), i.e. equal to the maximum number of electrons per f-sublevel (see paragraph 19 above). Element with Z= 72 (hafnium Hf) is an analogue of zirconium Zr and has been found in zirconium ores.

23. The next important conclusion from the analysis of table. 6 in paragraph 20 is the conclusion about the periodicity of filling the outer energy levels of atoms with electrons, which determines the periodicity of changes in the chemical properties of elements and their compounds.

Table 6

Electronic configurations of atoms
first 20 elements of the periodic table

Atomic number	Oboz- meaning	Layer	K	L	M	N
		n	1	2	3	4
		l	0	0, 1	0, 1, 2	0, 1, 2, 3
		Sublevel	1s	2s, 2p	3s, 3p, 3d	4s, 4p, 4d, 4f
		Number of electrons at a given sublevel
1 2	H He		1 2
3 4 5 6 7 8 9 10	Li Be B C N O F Ne		2 2 2 2 2 2 2 2	1, 0 2, 0 2, 1 2, 2 2, 3 2, 4 2, 5 2, 6
11 12 13 14 15 16 17 18	Na Mg Al Si P S Cl Ar		2 2 2 2 2 2 2 2	2, 6 2, 6 2, 6 2, 6 2, 6 2, 6 2, 6 2, 6	1, 0, 0 2, 0, 0 2, 1, 0 2, 2, 0 2, 3, 0 2, 4, 0 2, 5, 0 2, 6, 0
19 20	K Ca		2 2	2, 6 2, 6	2, 6, 0 2, 6, 0	1, 0, 0, 0 2, 0, 0, 0

Thus, the second period of D.I. Mendeleev’s table consists of eight elements with the following sublevels:

3Li 4 Be 5 B 6 C 7 N 8 O 9F 10 Ne 1s 2 2s 1 1s 2 2s 2 1s 2 2s 2 2p 1 1s 2 2s 2 2p 2 1s 2 2s 2 2p 3 1s 2 2s 2 2p 4 1s 2 2s 2 2p 5 1s 2 2s 2 2p 6

During the transition from lithium to neon, the charge of the atomic nucleus gradually increases from Z= 3 to Z= 10, which means that the forces of attraction of electrons to the nucleus increase, and as a result, the radii of the atoms of these elements decrease. Therefore, the ability of an atom to donate electrons (a typically metallic property), pronounced in the lithium atom, gradually weakens when moving from lithium to fluorine. The latter is a typical non-metal, that is, an element more capable than others of acquiring electrons.

Starting from the element next to neon (Na, Z= 11) the electronic structures of atoms are repeated, and therefore the electronic configurations of their outer electron shells are designated in a similar way ( n– period number):

ns 1 (Li,Na), ns 2 (Be, Mg), ns 2 n.p. 1 (B, Al), ns 2 n.p. 2 (C, Si) etc.

In the fourth period of D.I. Mendeleev’s table, transition elements appear that belong to secondary subgroups.

24. Elements belonging to the same subgroup have a similar arrangement of electrons in the outer electronic levels of the atoms. For example, the halogen atoms (the main subgroup of group VII) all have the electronic configuration ns 2 n.p. 5, and the atoms of elements of a side subgroup of the same group are characterized by an electronic configuration ( n– 1)s 2 (n– 1)p 6 (n– 1)d 5 ns 2 .

What is the essence of the similarities and differences between atoms of elements belonging to different subgroups of the same group of D.I. Mendeleev’s table? In the future, check your conclusions with Appendix 1 (P-24).

25. The numerical value of the valency of an atom, determined by the number of covalent chemical bonds formed by it, reflects the position of the element in D.I. Mendeleev’s PSCE. In many cases, the valency of an element atom in a compound is numerically equal to the group number in D.I. Mendeleev’s PSHE. However, there are exceptions to this rule. For example, the phosphorus atom on the outer (third, M) energy level contains three unpaired electrons (3 r-orbitals) and free valence cells d-orbitals. Consequently, the phosphorus atom is characterized by the so-called excitation electron, associated with the pairing of an electron pair and the transition of one of the resulting unpaired electrons to 3 d-orbital. For the excited state of the phosphorus atom, the formation of five covalent bonds is possible, and for the ground state - only three.

For the nitrogen atom, the excited state is atypical, since in this atom at the external energy level the number and state of electrons is the same as in the phosphorus atom, but there are no vacant cells, and only three electrons are missing for the completion and stability of this level.

Why then is the maximum valency of the nitrogen atom in compounds (i.e., the ability to form common electron pairs) not III, but IV?

26. Repeating paragraphs. 16, 17 of the methodological development, it is possible to explain the order of filling energy levels with electrons in the atoms of elements of the 4th large period of D.I. Mendeleev’s PSHE. The even series of this period begins with elements of the main subgroups - 39 K and 40 Ca, which are typical metals with constant valency, and already with element No. 21 ( Z= 21, Sc) then there are elements of side subgroups called d- elements or transitional. Try to explain the essence of these names and give relevant examples. In the future, check the correctness of your conclusions with Appendix 1 (P-26).

27. The chemical symbol of hydrogen H in D.I. Mendeleev’s PSHE is also placed in the main subgroup
Group I, and to the main subgroup of Group VII. Why is this acceptable? Check in the future the correctness of your conclusions in Appendix 1 (P-27).

When distributing electrons among quantum cells, the following guidelines follow:
Based on the Pauli principle: an atom cannot have two electrons with the same
set of values ​​of all quantum numbers, i.e., an atomic orbital cannot contain
press more than two electrons, and their spin moments should be opposite
opposite
↓

The notation system in general looks like this:

where p is the main one, ℓ is the orbital quantum number; x is the number of electrons,
in a given quantum state. For example, the 4d3 entry might be
interpreted as follows: three electrons occupy the fourth energy
Ski level, d-sublevel.

The nature of the development of energy sublevels determines the affiliation
element to one or another electronic family.

In s-elements, the external s-sublevel is built up, for example,

11 Na 1s2 2s2 2p6 3s1
In p-elements, the external p-sublevel is built up, for example,

9 F 1s 2s2 2p5 .

The s- and p-families include elements of the main subgroups of the periodic table.
tsy D.I. Mendeleev.

In d-elements, the d-sublevel of the penultimate level is built up,
For example,
2 2 6 2 6 2 2
22Ti 1s 2s 2p 3s 3p 3d 4s .

The d-family includes elements of side subgroups. The valency of this se-
families are s-electrons of the last energy level and d-electrons
penultimate level.

In the f-elements, the f-sublevel of the third external level is built,
For example,

58Се 1s22s22p63s23p63d l04s24p64d l04f l5s25p65d16s2.

Representatives of the f-electron family are lanthanides and actinides.

A quantum number can take two values: Therefore, no more than electrons can be in an atom in states with a given value:

Fundamentals of band theory

According to Bohr's postulates, in an isolated atom the energy of an electron can take strictly discrete values ​​(they also say that the electron is in one of the orbitals).

In the case of several atoms united by a chemical bond (for example, in a molecule), electron orbitals are split in an amount proportional to the number of atoms, forming so-called molecular orbitals. With a further increase in the system to a macroscopic crystal (the number of atoms is more than 10 20), the number of orbitals becomes very large, and the difference in the energies of electrons located in neighboring orbitals is correspondingly very small, the energy levels are split into almost continuous discrete sets - energy zones. The highest of the allowed energy bands in semiconductors and dielectrics, in which at a temperature of 0 K all energy states are occupied by electrons, is called the valence band, the next one is the conduction band. In metals, the conduction band is the highest allowed band in which electrons reside at a temperature of 0 K.

The band theory is based on the following main approximations:

1. The solid is a perfectly periodic crystal.

2. The equilibrium positions of the nodes of the crystal lattice are fixed, that is, the atomic nuclei are considered motionless (adiabatic approximation). Small vibrations of atoms around equilibrium positions, which can be described as phonons, are subsequently introduced as a perturbation of the electronic energy spectrum.

3. The many-electron problem is reduced to a single-electron one: the influence of all the others on a given electron is described by some averaged periodic field.

A number of essentially multielectron phenomena, such as ferromagnetism, superconductivity, and those where excitons play a role, cannot be consistently considered within the framework of band theory. At the same time, with a more general approach to constructing the theory of solids, it turned out that many results of the band theory are broader than its initial premises.

Photoconductivity.

Photoconductivity- the phenomenon of a change in the electrical conductivity of a substance upon absorption of electromagnetic radiation, such as visible, infrared, ultraviolet or x-ray radiation.

Photoconductivity is characteristic of semiconductors. The electrical conductivity of semiconductors is limited by the lack of charge carriers. When a photon is absorbed, an electron moves from the valence band to the conduction band. As a result, a pair of charge carriers is formed: an electron in the conduction band and a hole in the valence band. Both charge carriers, when voltage is applied to the semiconductor, create an electric current.

When photoconductivity is excited in an intrinsic semiconductor, the photon energy must exceed the band gap. In a doped semiconductor, the absorption of a photon can be accompanied by a transition from a level located in the bandgap, which allows the wavelength of light that causes photoconductivity to be increased. This circumstance is important for detecting infrared radiation. A condition for high photoconductivity is also a high light absorption rate, which is realized in direct-gap semiconductors

Quantum phenomena

37) Nuclear structure and radioactivity

Atomic nucleus- the central part of the atom, in which the bulk of its mass is concentrated (more than 99.9%). The nucleus is positively charged; the charge of the nucleus is determined by the chemical element to which the atom belongs. The sizes of the nuclei of various atoms are several femtometers, which is more than 10 thousand times smaller than the size of the atom itself.

The number of protons in the nucleus is called its charge number - this number is equal to the serial number of the element to which the atom belongs in Mendeleev’s table (Periodic Table of Elements). The number of protons in the nucleus determines the structure of the electron shell of a neutral atom and, thus, the chemical properties of the corresponding element. The number of neutrons in a nucleus is called its isotopic number. Nuclei with the same number of protons and different numbers of neutrons are called isotopes. Nuclei with the same number of neutrons but different numbers of protons are called isotones. The terms isotope and isotone are also used to refer to atoms containing these nuclei, as well as to characterize non-chemical varieties of a single chemical element. The total number of nucleons in a nucleus is called its mass number () and is approximately equal to the average mass of an atom shown in the periodic table. Nuclides with the same mass number but different proton-neutron composition are usually called isobars.

Radioactive decay(from lat. radius"beam" and āctīvus“effective”) - spontaneous change in composition (charge Z, mass number A) or the internal structure of unstable atomic nuclei by emission of elementary particles, gamma rays and/or nuclear fragments. The process of radioactive decay is also called radioactivity, and the corresponding nuclei (nuclides, isotopes and chemical elements) are radioactive. Substances containing radioactive nuclei are also called radioactive.