Written down 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 subscript at the top right indicates the number of electrons in the given sublevel. To compose the electronic formula of an atom of any element, it is enough to know the number of this element in the periodic table and 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. 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 shapes and properties of compounds of elements, are in periodic dependence of the magnitude of the atomic weights of the elements.
Modern formulation of the Periodic Law: the properties of elements, as well as the forms and properties of their compounds, are periodically dependent on the magnitude of the charge of the nucleus of their atoms.
Thus, the positive charge of the nucleus (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-external levels, can take part in the formation of chemical bonds. Such electrons are also called valence electrons.
Stoichiometric valence chemical 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 oxide SO 2 - 4, in oxide SO 3 -6.
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 state also 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 compounds has an oxidation state of -2; the exception is hydrogen peroxide Н 2 О 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, the oxidation state of oxygen in which 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 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 - 8. For example, the highest oxidation state of nitrogen (it is located in the fifth group) is +5 (in nitric acid and its salts), and the lowest is equal to -3 (in ammonia and ammonium salts).
7. The oxidation states of the elements in the compound compensate each other so that their sum for all atoms in a molecule or 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 towards more electronegative atoms (that is, assuming that the compound is composed only of ions).
The oxidation state corresponds to the number of electrons that must be attached to a positive ion in order to reduce it to a neutral atom, or subtract from a negative ion in order to oxidize it to a neutral atom:
Al 3+ + 3e - → Al
S 2− → S + 2e - (S 2− - 2e - → S)
The properties of elements, depending on the structure of the electron shell of the atom, vary by periods and groups of the periodic system. 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 of 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 = 0
K without transferring kinetic energy to the liberated 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 attaches to an atom in the gas phase at T = 0
K with the transformation of an atom into a negatively charged ion without transferring kinetic energy to the particle:
E + e- = E- + Ee.
Halogens, especially fluorine, have the maximum electron affinity (Ee = -328 kJ / mol).
The values of 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 ÷ = (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 decrease with an increase in the ordinal number of the element.
It should be emphasized that a constant value of electronegativity cannot be attributed to an element, 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 of the radii of atoms and ions in compounds, calculated from experimental data, are usually used. In this case, a distinction is made between covalent and metallic radii of atoms.
The dependence of the atomic and ionic radii on the charge of the nucleus of the atom of the element and is of a periodic nature... 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 abrupt, 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 the 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, a chemical bond is of an electronic nature, but it is carried out in different ways. Therefore, there are three main types of chemical bonds: covalent, ionic, metallic. 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. With an increase in the multiplicity, 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 its breaking, kJ / mol.
Covalent bond - A chemical bond formed by the sharing of a pair of electrons with two atoms.
The explanation of the chemical bond by the emergence 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, one more 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 the molecule are of electrical, Coulomb origin.
5. The bond is the 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:
Rice. 7. Exchange mechanism of covalent bond formation: a- 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, educated by donor-acceptor mechanism, refer to complex compounds
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Rice. 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.
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Directivity - t ... That is, the bond is formed in the direction of the maximum overlap of electron clouds . With respect to the line connecting the centers of the atoms forming the 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 the axis of the perpendicular line connecting the nuclei of the atom. The directionality of the bond determines the spatial structure of the molecules, that is, their geometric shape. Hybridization - it is a change in the shape of some orbitals during the formation of a covalent bond to achieve more effective overlapping of the orbitals. The chemical bond formed with the participation of the electrons of the hybrid orbitals is stronger than the bond with the participation of the electrons of the 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 °. The molecules in which sp-hybridization is carried out 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 °. The 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).
Rice. 10. Types of hybridizations of valence orbitals: a - sp-hybridization of valence orbitals; b - sp 2 - hybridization of valence orbitals; v - sp 3-hybridization of valence orbitals
Atom composition.
An atom consists of atomic nucleus and electronic shell.
The nucleus of an atom consists of protons ( p +) and neutrons ( n 0). Most hydrogen atoms have a single proton nucleus.
Number of protons N(p +) is equal to the nuclear charge ( Z) and the ordinal number of the element in the natural series of elements (and in the periodic table of elements).
N(p +) = Z
The sum of the number of neutrons N(n 0), denoted simply by the letter N, and the number of protons Z called massive number and denoted by the letter A.
A = Z + N
The electron shell of an atom consists of electrons moving around the nucleus ( e -).
Number of electrons N(e-) in the electron shell of a neutral atom is equal to the number of protons Z at its core.
The mass of a proton is approximately equal to the mass of a neutron and is 1840 times greater than the mass of an electron, so the mass of an atom is practically equal to the mass of a nucleus.
The shape of the atom is spherical. The radius of the nucleus is about 100,000 times smaller than the radius of the atom.
Chemical element- the kind of atoms (a set of atoms) with the same charge of the nucleus (with the same number of protons in the nucleus).
Isotope- a set of atoms of one element with the same number of neutrons in the nucleus (or the type of atoms with the same number of protons and the same number of neutrons in the nucleus).
Different isotopes differ from each other in the number of neutrons in the nuclei of their atoms.
The designation of a single atom or isotope: (E is the symbol of an element), for example:.
The structure of the electron shell of an atom
Atomic orbital- the state of an electron in an atom. Orbital symbol -. An electron cloud corresponds to each orbital.
Orbitals of real atoms in the ground (unexcited) state are of four types: s, p, d and f.
Electronic cloud- a part of space in which an electron can be detected with a probability of 90 (or more) percent.
Note: sometimes the concepts of "atomic orbital" and "electron cloud" are not distinguished, calling both the "atomic orbital".
The electron shell of an atom is layered. Electronic layer formed by electron clouds of the same size. Orbitals of one layer form electronic ("energy") level, their energies are the same for the hydrogen atom, but different for other atoms.
Similar orbitals of the same level are grouped into electronic (energy) sublevels:
s-sublevel (consists of one s-orbital), symbol -.
p-sublevel (consists of three p
d-sublevel (consists of five d-orbitals), symbol -.
f-sublevel (consists of seven f-orbitals), symbol -.
The energies of the orbitals of one sublevel are the same.
When designating sublevels, the number of the layer (electronic layer) is added to the symbol of the sublevel, for example: 2 s, 3p, 5d means s-sublevel of the second level, p-sublevel of the third level, d-sublevel of the fifth level.
The total number of sublevels in one level is equal to the number of the level n... The total number of orbitals at one level is n 2. Accordingly, the total number of clouds in one layer is also n 2 .
Designations: - free orbital (without electrons), - orbital with an unpaired electron, - orbital with an electron pair (with two electrons).
The order of filling the orbitals of an atom with electrons is determined by three laws of nature (formulations are given in a simplified manner):
1. The principle of least energy - electrons fill the orbitals in the order of increasing energy of the orbitals.
2. Pauli's principle - on one orbital there can be no more than two electrons.
3. Hund's rule - within the sublevel, electrons first fill free orbitals (one at a time), and only then form electron pairs.
The total number of electrons in the electronic level (or in the electronic layer) is 2 n 2 .
The distribution of sublevels by energy is expressed as follows (in the order of increasing energy):
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p ...
This sequence is clearly expressed in an energy diagram:
The distribution of electrons of an atom over levels, sublevels and orbitals (electronic configuration of an atom) can be depicted in the form of an electronic formula, an energy diagram, or, simply, in the form of a diagram of electronic layers ("electronic circuit").
Examples of the electronic structure of atoms:
Valence electrons- the electrons of the atom, which can take part in the formation of chemical bonds. For any atom, these are all external electrons plus those pre-external electrons, the energy of which is greater than that of the external ones. For example: a Ca atom has external electrons - 4 s 2, they are also valence; the Fe atom has external electrons - 4 s 2, but it has 3 d 6, therefore the iron atom has 8 valence electrons. The valence electronic formula of the calcium atom is 4 s 2, and the iron atom - 4 s 2 3d 6 .
Periodic table of chemical elements of D. I. Mendeleev
(natural system of chemical elements)
Periodic law of chemical elements(modern formulation): the properties of chemical elements, as well as simple and complex substances formed by them, are periodically dependent on the value of the charge from atomic nuclei.
Periodic system- graphic expression of the periodic law.
Natural range of chemical elements- a series of chemical elements, arranged according to the increasing number of protons in the nuclei of their atoms, or, which is the same, according to the increasing charges of the nuclei of these atoms. The ordinal number of an element in this row is equal to the number of protons in the nucleus of any atom of this element.
The table of chemical elements is constructed by "cutting" the natural series of chemical elements into periods(horizontal rows of the table) and groupings (vertical columns of the table) of elements with a similar electronic structure of atoms.
Depending on the method of combining elements into groups, the table may be long period(elements with the same number and type of valence electrons are collected in groups) and short-period(elements with the same number of valence electrons are collected in groups).
The groups of the short-period table are divided into subgroups ( the main and collateral) that coincide with the groups of the long-period table.
All atoms of elements of the same period have the same number of electronic layers, equal to the number of the period.
The number of elements in periods: 2, 8, 8, 18, 18, 32, 32. Most of the elements of the eighth period are obtained artificially, the last elements of this period have not yet been synthesized. All periods, except for the first one, begin with an element that forms an alkali metal (Li, Na, K, etc.), and end with an element that forms a noble gas (He, Ne, Ar, Kr, etc.).
In the short-period table there are eight groups, each of which is divided into two subgroups (main and secondary), in the long-period table there are sixteen groups, which are numbered in Roman numerals with the letters A or B, for example: IA, IIIB, VIA, VIIB. Group IA of the long-period table corresponds to the main subgroup of the first group of the short-period table; group VIIB - a side subgroup of the seventh group: the rest are similar.
The characteristics of chemical elements change naturally in groups and periods.
In periods (with an increase in the serial number)
- the charge of the nucleus increases,
- the number of external electrons increases,
- the radius of atoms decreases,
- the strength of the bond of electrons with the nucleus (ionization energy) increases,
- electronegativity increases,
- the oxidizing properties of simple substances are enhanced ("non-metallicity"),
- the reducing properties of simple substances ("metallicity") weaken,
- weakens the basic character of hydroxides and corresponding oxides,
- the acidic character of hydroxides and corresponding oxides increases.
In groups (with increasing serial number)
- the charge of the nucleus increases,
- the radius of atoms increases (only in A-groups),
- the bond strength of electrons with the nucleus decreases (ionization energy; only in A-groups),
- decreases electronegativity (only in A-groups),
- the oxidizing properties of simple substances weaken ("non-metallic"; only in A-groups),
- the reducing properties of simple substances are enhanced ("metallicity"; only in A-groups),
- the basic character of hydroxides and corresponding oxides increases (only in A-groups),
- the acidic nature of hydroxides and corresponding oxides weakens (only in A-groups),
- the stability of hydrogen compounds decreases (their reductive activity increases; only in A-groups).
Tasks and tests on the topic "Topic 9." The structure of the atom. DI Mendeleev's Periodic Law and Periodic Table of Chemical Elements (PSKhE) "."
- Periodic law - Periodic law and the structure of atoms 8-9 grade
You should know: the laws of filling orbitals with electrons (the principle of least energy, Pauli's principle, Hund's rule), the structure of the periodic table of elements.You should be able to: determine the composition of an atom by the position of an element in the periodic system, and, conversely, find an element in the periodic system, knowing its composition; depict the structure diagram, the electronic configuration of an atom, ion, and, conversely, determine the position of a chemical element in the PSCE according to the diagram and electronic configuration; to characterize the element and the substances formed by it according to its position in the PSCE; determine changes in the radius of atoms, properties of chemical elements and the substances formed by them within one period and one main subgroup of the periodic system.
Example 1. Determine the number of orbitals at the third electronic level. What are these orbitals?
To determine the number of orbitals, we use the formula N orbitals = n 2, where n- level number. N orbitals = 3 2 = 9. One 3 s-, three 3 p- and five 3 d-orbitals.Example 2. Determine which atom of which element has electronic formula 1 s 2 2s 2 2p 6 3s 2 3p 1 .
In order to determine which element it is, it is necessary to find out its serial number, which is equal to the total number of electrons of the atom. In this case: 2 + 2 + 6 + 2 + 1 = 13. This is aluminum.After making sure that everything you need is learned, proceed to the tasks. We wish you every success.
Recommended reading:- OS Gabrielyan and others. Chemistry 11 class. M., Bustard, 2002;
- G.E. Rudzitis, F.G. Feldman. Chemistry 11 cl. M., Education, 2001.
6.6. Features of the electronic structure of atoms of chromium, copper and some other elements
If you carefully looked at Appendix 4, you probably noticed that the sequence of filling the orbitals with electrons is violated in the atoms of some elements. Sometimes these violations are called "exceptions", but they are not - there are no exceptions to the laws of Nature!
The first element with this violation is chrome. Let's consider in more detail its electronic structure (Fig. 6.16 a). The chromium atom has 4 s-sub-level not two, as one would expect, but only one electron. But at 3 d-sublevel five electrons, but this sublevel is filled after 4 s-sublevel (see Fig. 6.4). To understand why this is happening, let's see what electron clouds are 3 d is the sublevel of this atom.
Each of five 3 d-clouds in this case is formed by one electron. As you already know from § 4 of this chapter, the common electron cloud of these five electrons has a spherical shape, or, as they say, spherically symmetric. By the nature of the electron density distribution in different directions, it is similar to 1 s-EO. The energy of the sublevel, the electrons of which form such a cloud, turns out to be less than in the case of a less symmetric cloud. In this case, the energy of the orbitals is 3 d-sublevel is equal to energy 4 s-orbital. When symmetry is broken, for example, when the sixth electron appears, the energy of the orbitals is 3 d-sublevel again becomes larger than energy 4 s-orbital. Therefore, the manganese atom again has a second electron by 4 s-AO.
The general cloud of any sublevel, filled with electrons, both half and completely, has spherical symmetry. The decrease in energy in these cases is of a general nature and does not depend on whether any sublevel is half or completely filled with electrons. And if so, then we should look for the next violation in the atom, into the electronic shell of which the ninth "comes" last d-electron. Indeed, the copper atom has 3 d-sublayer 10 electrons, and on 4 s- there is only one sublevel (Fig. 6.16 b).
The decrease in the energy of the orbitals of a fully or half-filled sublevel is the cause of a number of important chemical phenomena, some of which you will become acquainted with.
6.7. Outer and valence electrons, orbitals and sublevels
In chemistry, the properties of isolated atoms, as a rule, are not studied, since almost all atoms, being part of various substances, form chemical bonds. Chemical bonds are formed when the electronic shells of atoms interact. For all atoms (except for hydrogen), not all electrons take part in the formation of chemical bonds: boron has three electrons out of five, carbon has four out of six, and, for example, barium has two out of fifty-six. These "active" electrons are called valence electrons.
Sometimes valence electrons are confused with external electrons, and they are not the same thing.
The electron clouds of the outer electrons have the maximum radius (and the maximum value of the principal quantum number).
It is the outer electrons that take part in the formation of a bond in the first place, if only because when atoms approach each other, the electron clouds formed by these electrons come into contact first of all. But together with them, part of the electrons can take part in the formation of a bond. pre-external(penultimate) layer, but only if they have an energy that does not differ much from the energy of the outer electrons. Both those and other electrons of an atom are valence. (In lanthanides and actinides, even some "pre-external" electrons are valence)
The energy of valence electrons is much higher than the energy of other electrons of the atom, and valence electrons differ significantly less in energy from each other.
External electrons are always valence only if the atom can form chemical bonds at all. So, both electrons of the helium atom are external, but they cannot be called valence, since the helium atom does not form any chemical bonds at all.
Valence electrons occupy valence orbitals, which in turn form valence sublevels.
As an example, consider an iron atom, the electronic configuration of which is shown in Fig. 6.17. Of the electrons of the iron atom, the maximum principal quantum number ( n= 4) have only two 4 s-electron. Therefore, it is they who are the outer electrons of this atom. The outer orbitals of the iron atom are all orbitals with n= 4, and the outer sublevels are all sublevels formed by these orbitals, that is, 4 s-,
4p-, 4d- and 4 f-EPU.
The outer electrons are always valence, therefore 4 s-electrons of the iron atom - valence electrons. And if so, then 3 d-electrons with slightly higher energy will also be valence. At the outer level of the iron atom, in addition to the filled 4 s-AO there are still free 4 p-, 4d- and 4 f-AO. They are all external, but there are only 4 valences among them. R-AO, since the energy of the other orbitals is much higher, and the appearance of electrons in these orbitals is not beneficial for the iron atom.
So the iron atom
external electronic level - fourth,
external sublevels - 4 s-, 4p-, 4d- and 4 f-EPU,
outer orbitals - 4 s-, 4p-, 4d- and 4 f-AO,
outer electrons - two 4 s-electron (4 s 2),
outer electron layer - fourth,
external electronic cloud - 4 s-EO
valence sublevels - 4 s-, 4p-, and 3 d-EPU,
valence orbitals - 4 s-, 4p-, and 3 d-AO,
valence electrons - two 4 s-electron (4 s 2) and six 3 d-electrons (3 d 6).
The valence sublevels can be partially or completely filled with electrons, or they can generally remain free. With an increase in the nuclear charge, the values of the energy of all sublevels decrease, but due to the interaction of electrons with each other, the energy of different sublevels decreases with different "rates". The energy is completely filled d- and f-sublevels decreases so much that they cease to be valence.
As an example, consider the atoms of titanium and arsenic (Fig. 6.18).
In the case of the titanium atom 3 d-The EPU is only partially filled with electrons, and its energy is greater than the energy 4 s-EPU, and 3 d-electrons are valence. Arsenic atom has 3 d-The EPU is completely filled with electrons, and its energy is significantly less than the energy 4 s-EPU, and therefore 3 d-electrons are not valence.
In the examples given, we have analyzed valence electronic configuration atoms of titanium and arsenic.
The valence electronic configuration of an atom is depicted as valence electronic formula, or in the form energy diagram of valence sublevels.
VALENT ELECTRONS, EXTERNAL ELECTRONS, VALENT EPU, VALENT AO, VALENT ELECTRONIC CONFIGURATION OF THE ATOM, VALENCE ELECTRONIC FORMULA, VALENCE SUB-LEVEL DIAGRAM.
1. On the energy diagrams you have drawn up and in the complete electronic formulas of the atoms Na, Mg, Al, Si, P, S, Cl, Ar, indicate the outer and valence electrons. Make up the valence electronic formulas of these atoms. On the energy diagrams, highlight the parts corresponding to the energy diagrams of the valence sublevels.
2. What is common between the electronic configurations of atoms a) Li and Na, B and Al, O and S, Ne and Ar; b) Zn and Mg, Sc and Al, Cr and S, Ti and Si; c) H and He, Li and O, K and Kr, Sc and Ga. What are their differences
3.How many valence sublevels in the electron shell of an atom of each of the elements: a) hydrogen, helium and lithium, b) nitrogen, sodium and sulfur, c) potassium, cobalt and germanium
4. How many valence orbitals are completely filled in the atom of a) boron, b) fluorine, c) sodium?
5.How many orbitals with an unpaired electron in an atom of a) boron, b) fluorine, c) iron
6. How many free outer orbitals does a manganese atom have? And how many free valences?
7. For the next lesson, prepare a strip of paper 20 mm wide, divide it into cells (20 × 20 mm), and apply a natural range of elements (from hydrogen to meitnerium) to this strip.
8.In each cell, place the symbol of the element, its ordinal number and the valence electronic formula, as shown in Fig. 6.19 (use Appendix 4).
6.8. Systematization of atoms according to the structure of their electronic shells
The systematization of chemical elements is based on the natural series of elements
and principle of similarity of electronic shells their atoms.
You are already familiar with the natural range of chemical elements. Now let's get acquainted with the principle of similarity of electronic shells.
Considering the valence electronic formulas of atoms in the NRE, it is easy to find that for some atoms they differ only in the values of the principal quantum number. For example 1 s 1 for hydrogen, 2 s 1 for lithium, 3 s 1 for sodium, etc., or 2 s 2 2p 5 for fluorine, 3 s 2 3p 5 for chlorine, 4 s 2 4p 5 for bromine, etc. This means that the outer regions of the clouds of valence electrons of such atoms are very similar in shape and differ only in size (and, of course, in electron density). And if so, then the electron clouds of such atoms and the corresponding valence configurations can be called like... For atoms of different elements with similar electronic configurations, we can write general valence electronic formulas: ns 1 in the first case and ns 2 np 5 in the second. Moving along the natural series of elements, one can find other groups of atoms with similar valence configurations.
Thus, atoms with similar valence electronic configurations are regularly found in the natural series of elements.
This is the principle of similarity of electronic shells.
Let's try to identify the kind of this regularity. To do this, we will use the natural row of elements you have made.
ERE begins with hydrogen, the valence electronic formula of which is 1 s 1 . In search of similar valence configurations, we cut the natural series of elements before elements with a common valence electronic formula ns 1 (i.e., before lithium, before sodium, etc.). We got the so-called "periods" of the elements. Let's add the resulting "periods" so that they become the rows of the table (see Fig. 6.20). As a result, only the atoms of the first two columns of the table will have similar electronic configurations.
Let's try to achieve similarity of valence electronic configurations in other columns of the table. To do this, we cut out from the 6th and 7th periods elements with numbers 58 - 71 and 90 - 103 (they are filled with 4 f- and 5 f-sub-levels) and place them below the table. Let's move the symbols of the remaining elements horizontally as shown in the figure. After that, atoms of elements standing in one column of the table will have similar valence configurations, which can be expressed by general valence electronic formulas: ns 1 , ns 2 , ns 2 (n–1)d 1 , ns 2 (n–1)d 2 and so on until ns 2 np 6. All deviations from the general valence formulas are due to the same reasons as in the case of chromium and copper (see paragraph 6.6).
As you can see, using the ERE and applying the principle of the similarity of electronic shells, we were able to systematize the chemical elements. Such a system of chemical elements is called natural, since it is based solely on the laws of Nature. The table we received (Fig. 6.21) is one of the ways to graphically depict the natural system of elements and is called long-period table of chemical elements.
PRINCIPLE OF SIMILARITY OF ELECTRONIC SHELLS, NATURAL SYSTEM OF CHEMICAL ELEMENTS ("PERIODIC" SYSTEM), TABLE OF CHEMICAL ELEMENTS.
6.9. Long-period table of chemical elements
Let's take a closer look at the structure of the long-period table of chemical elements.
The rows in this table, as you already know, are called "periods" of the elements. The periods are numbered with Arabic numerals from 1 to 7. There are only two elements in the first period. The second and third periods, each containing eight elements, are called short periods. The fourth and fifth periods, each containing 18 elements, are called long periods. The sixth and seventh periods, each containing 32 elements, are called extra-long periods.
The columns in this table are named in groups elements. Group numbers are designated by Roman numerals with Latin letters A or B.
Elements of some groups have their own common (group) names: elements of the IA group (Li, Na, K, Rb, Cs, Fr) - alkaline elements(or alkali metal elements); Group IIA elements (Ca, Sr, Ba and Ra) - alkaline earth elements(or alkaline earth metal elements) (the name "alkali metals" and alkaline earth metals "refers to simple substances formed by the corresponding elements and should not be used as names for groups of elements); elements of group VIA (O, S, Se, Te, Po) - chalcogenes, elements of VIIA group (F, Cl, Br, I, At) - halogens, elements of VIIIA group (He, Ne, Ar, Kr, Xe, Rn) - noble gas elements. (The traditional name "noble gases" also refers to simple substances)
Elements with serial numbers 58 - 71 (Ce - Lu) usually taken out to the bottom of the table are called lanthanides("following lanthanum"), and elements with serial numbers 90 - 103 (Th - Lr) - actinides("following anemones"). There is a variant of the long-period table, in which lanthanides and actinides are not cut out from the NRE, but remain in place in super-long periods. This table is sometimes called superlong-period.
Long period table is divided into four block(or section).
s-Block includes elements of IA and IIA-groups with common valence electronic formulas ns 1 and ns 2
(s-elements).
r-Block includes elements from IIIA to VIIIA group with common valence electronic formulas from ns 2 np 1 to ns 2 np 6 (p-elements).
d-block includes elements from IIIB to IIB group with general valence electronic formulas from ns 2 (n–1)d 1 to ns 2 (n–1)d 10 (d-elements).
f-Block includes lanthanides and actinides ( f-elements).
The elements s- and p-blocks form A-groups, and the elements d-block - B-group of the system of chemical elements. Everything f-elements are formally included in the IIIB group.
The elements of the first period - hydrogen and helium - are s-elements and can be placed in groups IA and IIA. But helium is more often placed in the VIIIA group as an element that ends the period, which fully corresponds to its properties (helium, like all other simple substances formed by the elements of this group, is a noble gas). Hydrogen, on the other hand, is often placed in the VIIA group, since in its properties it is much closer to halogens than to alkaline elements.
Each of the periods of the system begins with an element with a valence configuration of atoms ns 1, since it is from these atoms that the formation of the next electron layer begins, and ends with an element with a valence configuration of atoms ns 2 np 6 (except for the first period). This makes it easy to distinguish groups of sublevels on the energy diagram, which are filled with electrons from atoms of each of the periods (Fig. 6.22). Do this work with all the sublevels shown in your copy of Figure 6.4. The sublevels highlighted in Figure 6.22 (except for completely filled d- and f-sublevels) are valence for atoms of all elements of a given period.
Appearance in periods s-, p-, d- or f-elements fully correspond to the filling sequence s-, p-, d- or f-sub-level electrons. This feature of the system of elements allows, knowing the period and the group that this element belongs to, immediately write down its valence electronic formula.
LONG PERIOD TABLE OF CHEMICAL ELEMENTS, BLOCKS, PERIODS, GROUPS, ALKALINE ELEMENTS, ALKALINE EARTH ELEMENTS, CHALCOGENS, HALOGENS, ELEMENTS OF NOBLE GASES, LATHANOIDS.
Write down the general valence electronic formulas of atoms of elements a) IVA and IVB groups, b) IIIA and VIIB groups?
2. What is common between the electronic configurations of atoms of elements A and B groups? How do they differ?
3. How many groups of elements are included in a) s-block, b) R-block, c) d-block?
4. Continue Figure 30 towards increasing the energy of the sublevels and select the groups of sublevels that are filled with electrons in the 4th, 5th and 6th periods.
5. List the valence sublevels of a) calcium, b) phosphorus, c) titanium, d) chlorine, e) sodium atoms. 6.Formulate how s-, p- and d-elements differ from each other.
7. Explain why the belonging of an atom to any element is determined by the number of protons in the nucleus, and not by the mass of this atom.
8. For atoms of lithium, aluminum, strontium, selenium, iron and lead, draw up valence, complete and abbreviated electronic formulas and draw energy diagrams of valence sublevels. 9. Atoms of which elements correspond to the following valence electronic formulas: 3 s 1 , 4s 1 3d 1, 2s 2 2 p 6 , 5s 2 5p 2 , 5s 2 4d 2 ?
6.10. Types of electronic formulas of the atom. Algorithm for their compilation
For various purposes, we need to know either the full or the valence configuration of the atom. Each of these electronic configurations can be represented by both a formula and an energy diagram. That is, complete electronic configuration of an atom expressed the complete electronic formula of the atom, or complete energy diagram of the atom... In turn, valence electron configuration of an atom expressed valence(or, as it is often called, " short ") electronic formula of the atom, or diagram of the valence sublevels of the atom(fig. 6.23).
We used to make the electronic formulas of atoms using the ordinal numbers of the elements. In this case, we determined the sequence of filling the sublevels with electrons according to the energy diagram: 1 s, 2s,
2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p,
6s, 4f, 5d, 6p, 7s etc. And only by writing down the complete electronic formula, we could write down the valence formula.
The valence electronic formula of the atom, which is most often used, is more convenient to write based on the position of the element in the system of chemical elements, according to the coordinates period - group.
Let's take a closer look at how this is done for elements. s-, p- and d-blocks.
For items s-block valence electronic formula of an atom consists of three symbols. In general, it can be written as follows:
In the first place (in place of the large cell), the period number is put (equal to the main quantum number of these s-electrons), and on the third (in the superscript) is the group number (equal to the number of valence electrons). Taking magnesium atom as an example (3rd period, IIA group), we get:
For items p-block valence electronic formula of an atom consists of six symbols:
Here, instead of large cells, the period number is also put (equal to the main quantum number of these s- and p-electrons), and the group number (equal to the number of valence electrons) turns out to be equal to the sum of the superscripts. For the oxygen atom (2nd period, group VIA) we get:
2s 2 2p 4 .
Valence electronic formula of most elements d-block can be written like this:
As in the previous cases, here, instead of the first cell, the period number is put (equal to the main quantum number of these s-electrons). The number in the second cell turns out to be one less, since the main quantum number of these d-electrons. The group number here is also equal to the sum of the indices. Example - the valence electronic formula of titanium (4th period, IVB group): 4 s 2 3d 2 .
The group number is equal to the sum of the indices and for the elements of the VIB group, but they, as you remember, on the valence s- there is only one electron sublevel, and the general valence electronic formula ns 1 (n–1)d 5 . Therefore, the valence electronic formula, for example, molybdenum (5th period) - 5 s 1 4d 5 .
It is just as easy to compose the valence electronic formula of any element of the IB group, for example, gold (6th period)> -> 6 s 1 5d 10, but in this case it must be remembered that d- the electrons of the atoms of the elements of this group still remain valence, and some of them can participate in the formation of chemical bonds.
The general valence electronic formula of atoms of group IIB elements is ns 2 (n – 1)d ten . Therefore, the valence electronic formula, for example, of a zinc atom is 4 s 2 3d 10 .
The valence electronic formulas of the elements of the first triad (Fe, Co and Ni) also obey the general rules. Iron, an element of group VIIIB, has a valence electronic formula of 4 s 2 3d 6. The cobalt atom has one d-electron is larger (4 s 2 3d 7), and for the nickel atom - by two (4 s 2 3d 8).
Using only these rules for writing valence electronic formulas, it is impossible to compose the electronic formulas of atoms of some d-elements (Nb, Ru, Rh, Pd, Ir, Pt), since their filling of valence sublevels with electrons due to the tendency to highly symmetric electron shells has some additional features.
Knowing the valence electronic formula, it is possible to write down the complete electronic formula of the atom (see below).
Often, instead of cumbersome full electronic formulas, one writes abbreviated electronic formulas atoms. To compose them in the electronic formula, all the electrons of the atom except for the valence ones are selected, their symbols are placed in square brackets and the part of the electronic formula corresponding to the electronic formula of the atom of the last element of the previous period (the element that forms the noble gas) is replaced with the symbol of this atom.
Examples of electronic formulas of different types are shown in Table 14.
Table 14. Examples of electronic formulas of atoms
Electronic formulas |
|||
Abbreviated |
Valence |
||
1s 2 2s 2 2p 3 |
2s 2 2p 3 |
2s 2 2p 3 |
|
1s 2 2s 2 2p 6 3s 2 3p 5 |
3s 2 3p 5 |
3s 2 3p 5 |
|
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 5 |
4s 2 3d 5 |
4s 2 3d 5 |
|
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 3 |
4s 2 4p 3 |
4s 2 4p 3 |
|
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 |
4s 2 4p 6 |
4s 2 4p 6 |
|
Algorithm for drawing up the electronic formulas of atoms (for example, the iodine atom)
№ |
Operation |
Result |
|
Determine the coordinates of the atom in the table of elements. |
Period 5, group VIIA |
||
Make a valence electronic formula. |
5s 2 5p 5 |
||
Complete the symbols of the internal electrons in the sequence of filling the sublevels with them. |
1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 5 |
||
Considering the decrease in energy of fully filled d- and f-Sub-levels, write down the complete electronic formula. |
|
||
Note the valence electrons. |
1s 2 2s 2 2p 6 3s 2 3p 6 3d 10 4s 2 4p 6 4d 10 5s 2 5p 5 |
||
Highlight the electronic configuration of the preceding noble gas atom. |
|||
Write down the abbreviated electronic formula, combining all in square brackets non-bonded electrons. |
5s 2 5p 5 |
Notes (edit)
1. For elements of the 2nd and 3rd periods, the third operation (without the fourth) immediately leads to a complete electronic formula.
2. (n – 1)d 10 -Electrons remain valence at the atoms of the elements of group IB.
COMPLETE ELECTRONIC FORMULA, VALENCE ELECTRONIC FORMULA, REDUCED ELECTRONIC FORMULA, ALGORITHM FOR COMPOSITION OF ELECTRONIC FORMULAS OF ATOMS.
1. Make the valence electronic formula of the atom of element a) the second period of the third A group, b) the third period of the second A group, c) the fourth period of the fourth A group.
2. Make the abbreviated electronic formulas of the atoms of magnesium, phosphorus, potassium, iron, bromine and argon.
6.11. Short-period table of chemical elements
For more than 100 years that have passed since the discovery of the natural system of elements, several hundred of the most diverse tables have been proposed, graphically reflecting this system. Of these, in addition to the long-period table, the most widespread is the so-called short-period table of elements by D.I.Mendeleev. A short-period table is obtained from a long-period table if the 4th, 5th, 6th and 7th periods are cut before the elements of the IB group, move apart and fold the resulting rows as we used to add the periods. The result is shown in Figure 6.24.
Lanthanides and actinides are also placed under the main table here.
V groups This table contains elements whose atoms equal number of valence electrons no matter what orbitals these electrons are in. So, the elements are chlorine (a typical element that forms a non-metal; 3 s 2 3p 5) and manganese (metal-forming element; 4 s 2 3d 5), not possessing the semblance of electronic shells, here fall into the same seventh group. The need to distinguish between such elements forces you to select in groups subgroups: the main- analogs of the A-groups of the long-period table and collateral- analogs of B-groups. In Figure 34, the symbols of the elements of the main subgroups are shifted to the left, and the elements of the secondary subgroups are shifted to the right.
True, this arrangement of elements in the table also has its advantages, because it is the number of valence electrons that primarily determines the valence capabilities of an atom.
The long-period table reflects the regularities of the electronic structure of atoms, the similarity and regularities of changes in the properties of simple substances and compounds by groups of elements, the regular change in a number of physical quantities characterizing atoms, simple substances and compounds throughout the entire system of elements, and much more. The short-period table is less convenient in this respect.
SHORT PERIOD TABLE, MAIN SUBGROUPS, SIDE SUBGROUPS.
1.Convert the long-period table you built from a natural series of elements into a short-period one. Reverse the transformation.
2. Is it possible to draw up a general valence electronic formula of atoms of elements of one group of a short-period table? Why?
6.12. Sizes of atoms. Orbital radii
.The atom has no clear boundaries. What is considered the size of an isolated atom? The nucleus of an atom is surrounded by an electron shell, and the shell consists of electron clouds. The size of the EO is characterized by the radius r eo. All clouds in the outer layer have approximately the same radius. Therefore, the size of an atom can be characterized by this radius. It is called orbital radius of the atom(r 0).
The values of the orbital radii of the atoms are given in Appendix 5.
The radius of the EO depends on the charge of the nucleus and on which orbital the electron that forms this cloud is located in. Consequently, the orbital radius of an atom depends on the same characteristics.
Consider the electron shells of hydrogen and helium atoms. Both in the hydrogen atom and in the helium atom, the electrons are at 1 s-AO, and their clouds would have the same size if the charges of the nuclei of these atoms were the same. But the charge of the nucleus of a helium atom is twice as large as the charge of the nucleus of a hydrogen atom. According to Coulomb's law, the force of attraction acting on each of the electrons of a helium atom is twice the force of attraction of an electron to the nucleus of a hydrogen atom. Consequently, the radius of the helium atom must be much smaller than the radius of the hydrogen atom. This is true: r 0 (He) / r 0 (H) = 0.291 E / 0.529 E 0.55.
The lithium atom has an outer electron at 2 s-AO, that is, it forms a cloud of the second layer. Naturally, its radius should be larger. Really: r 0 (Li) = 1.586 E.
The atoms of the remaining elements of the second period have external electrons (and 2 s, and 2 p) are located in the same second electron layer, and the nuclear charge of these atoms increases with increasing serial number. Electrons are more strongly attracted to the nucleus, and, naturally, the radii of the atoms decrease. We could repeat this reasoning for atoms of elements of other periods, but with one clarification: the orbital radius decreases monotonically only when each of the sublevels is filled.
But if we ignore the particulars, then the general nature of the change in the size of atoms in the system of elements is as follows: with an increase in the ordinal number in the period, the orbital radii of the atoms decrease, and in the group, they increase. The largest atom is a cesium atom, and the smallest is a helium atom, but of the atoms of elements that form chemical compounds (helium and neon do not form them), the smallest is a fluorine atom.
Most of the atoms of the elements that are in the natural row after the lanthanides have orbital radii somewhat smaller than one would expect based on general laws. This is due to the fact that 14 lanthanides are located between lanthanum and hafnium in the system of elements, and, therefore, the charge of the nucleus of the hafnium atom is 14 e more than lanthanum. Therefore, the outer electrons of these atoms are attracted to the nucleus more strongly than they would in the absence of lanthanides (this effect is often called "lanthanide compression").
Note that when going from the atoms of the elements of the VIIIA group to the atoms of the elements of the IA group, the orbital radius increases abruptly. Consequently, our choice of the first elements of each period (see § 7) turned out to be correct.
ORBITAL RADIUS OF THE ATOM, ITS CHANGE IN THE SYSTEM OF ELEMENTS.
1.According to the data given in Appendix 5, plot on graph paper a graph of the dependence of the orbital radius of an atom on the ordinal number of an element for elements with Z from 1 to 40. The length of the horizontal axis is 200 mm, the length of the vertical axis is 100 mm.
2. How can you characterize the appearance of the resulting broken line?
6.13. Ionization energy of an atom
If you impart additional energy to an electron in an atom (how this can be done, you will learn from a physics course), then the electron can go to another AO, that is, the atom will be in excited state... This state is unstable, and the electron will almost immediately return to its original state, and the excess energy will be released. But if the energy imparted to the electron is large enough, the electron can completely detach from the atom, while the atom ionizes, that is, it turns into a positively charged ion ( cation). The energy required for this is called the ionization energy of the atom(E and).
It is rather difficult to tear an electron from a single atom and measure the energy required for this, therefore, it is practically determined and used molar ionization energy(E and m).
The molar ionization energy shows what is the smallest energy that is required to detach 1 mole of electrons from 1 mole of atoms (one electron from each atom). This value is usually measured in kilojoules per mole. The values of the molar ionization energy of the first electron for most elements are given in Appendix 6.
How does the ionization energy of an atom depend on the position of an element in a system of elements, that is, how does it change in a group and a period?
According to the physical meaning, the ionization energy is equal to the work that must be spent on overcoming the force of attraction of the electron to the atom when the electron moves from the atom to an infinite distance from it.
where q- electron charge, Q Is the charge of the cation remaining after the removal of the electron, and r o is the orbital radius of the atom.
AND q, and Q- the quantities are constant, and we can conclude that the work on the separation of an electron A, and with it the ionization energy E and, are inversely proportional to the orbital radius of the atom.
Having analyzed the values of the orbital radii of atoms of various elements and the corresponding values of the ionization energy given in Appendices 5 and 6, you can make sure that the relationship between these values is close to proportional, but somewhat different from it. The reason that our conclusion does not agree very well with the experimental data is that we used a very crude model that does not take into account many significant factors. But even this rough model allowed us to draw the correct conclusion that with an increase in the orbital radius, the ionization energy of an atom decreases and, conversely, with a decrease in the radius, it increases.
Since the orbital radius of the atoms decreases in the period with an increase in the ordinal number, the ionization energy increases. In a group, with an increase in the ordinal number, the orbital radius of the atoms, as a rule, increases, and the ionization energy decreases. The largest molar ionization energy is found in the smallest atoms, helium atoms (2372 kJ / mol), and among the atoms capable of forming chemical bonds, in fluorine atoms (1681 kJ / mol). The smallest is for the largest atoms, cesium atoms (376 kJ / mol). In a system of elements, the direction of increasing the ionization energy can be schematically shown as follows: 
In chemistry, it is important that the ionization energy characterizes the tendency of an atom to give up "its" electrons: the greater the ionization energy, the less the atom is inclined to give up electrons, and vice versa.
EXCITED STATE, IONIZATION, CATION, IONIZATION ENERGY, IONIZATION MOLAR ENERGY, IONIZATION ENERGY CHANGE IN THE ELEMENT SYSTEM.
1.Using the data given in Appendix 6, determine how much energy needs to be spent in order to take one electron away from all sodium atoms with a total mass of 1 g.
2. Using the data given in Appendix 6, determine how many times more energy needs to be spent to detach one electron from all sodium atoms with a mass of 3 g than from all potassium atoms of the same mass. Why is this ratio different from the ratio of the molar ionization energies of the same atoms?
3.According to the data given in Appendix 6, build a graph of the dependence of the molar ionization energy on the serial number for elements with Z from 1 to 40. The dimensions of the graph are the same as in the task to the previous paragraph. See if this schedule matches the selection of "periods" of the system of elements.
6.14. Electron affinity energy
.The second most important energy characteristic of an atom is electron affinity energy(E with).
In practice, as in the case of the ionization energy, the corresponding molar quantity is usually used - molar electron affinity energy().
The molar energy of affinity for an electron shows what is the energy released when one mole of electrons is attached to one mole of neutral atoms (one electron to each atom). Like the molar ionization energy, this value is also measured in kilojoules per mole.
At first glance, it may seem that energy should not be released in this case, because an atom is a neutral particle, and there are no electrostatic forces of attraction between a neutral atom and a negatively charged electron. On the contrary, approaching an atom, an electron, it would seem, should repel from the same negatively charged electrons that form an electron shell. Actually this is not true. Remember if you have ever had to deal with atomic chlorine. Of course not. After all, it only exists at very high temperatures. Even the more stable molecular chlorine is practically not found in nature - if necessary, it has to be obtained using chemical reactions. And with sodium chloride (table salt) you have to deal with it all the time. After all, table salt is consumed every day by a person with food. And in nature, it occurs quite often. But the composition of table salt includes chloride ions, that is, chlorine atoms, which have attached one "extra" electron. One of the reasons for this prevalence of chloride ions is that chlorine atoms have a tendency to attach electrons, that is, when chloride ions are formed from chlorine atoms and electrons, energy is released.
One of the reasons for the release of energy is already known to you - it is associated with an increase in the symmetry of the electron shell of the chlorine atom during the transition to a singly charged anion... At the same time, as you remember, energy 3 p-sublevel decreases. There are other more complex reasons as well.
Due to the fact that the value of the electron affinity energy is influenced by several factors, the nature of the change in this value in the system of elements is much more complex than the nature of the change in the ionization energy. You can be convinced of this by analyzing the table given in Appendix 7. But since the value of this quantity is determined, first of all, by the same electrostatic interaction as the values of the ionization energy, then its change in the system of elements (at least in A- groups), in general terms, is similar to a change in the ionization energy, that is, the energy of affinity for an electron in a group decreases, and in a period it increases. It is maximal for fluorine (328 kJ / mol) and chlorine (349 kJ / mol) atoms. The nature of the change in the energy of affinity for an electron in a system of elements resembles the nature of the change in the ionization energy, that is, the direction of an increase in the energy of affinity for an electron can be schematically shown as follows:
2. On the same scale along the horizontal axis as in the previous tasks, plot the dependence of the molar energy of affinity for an electron on the ordinal number for atoms of elements with Z from 1 to 40 using app 7.
3.What is the physical meaning of negative values of the electron affinity energy?
4. Why, of all atoms of the elements of the 2nd period, only beryllium, nitrogen and neon have negative values of the molar energy of affinity for an electron?
6.15. The tendency of atoms to give and attach electrons
You already know that the propensity of an atom to give up its own and attach foreign electrons depends on its energy characteristics (ionization energy and electron affinity energy). Which atoms are more inclined to donate their electrons, and which ones are more inclined to accept others?
To answer this question, let us summarize in Table 15 everything that we know about the change in these tendencies in the system of elements.
Table 15. Change in the propensity of atoms to give up their own and attach foreign electrons
- First, we write down the chemical sign. element, where below to the left of the sign we indicate its serial number.
- Further, by the number of the period (from which the element), we determine the number of energy levels and draw such a number of arcs next to the sign of the chemical element.
- Then, according to the group number, the number of electrons at the outer level, we write down under the arc.
- At the 1st level the maximum possible is 2e, at the second it is already 8, at the third - as many as 18. We begin to put numbers under the corresponding arcs.
- The number of electrons at the penultimate level must be calculated as follows: the number of electrons already affixed is subtracted from the ordinal number of the element.
- It remains to turn our circuit into an electronic formula:
- We write a chemical element and its serial number. The number shows the number of electrons in an atom.
- We draw up a formula. To do this, you need to find out the number of energy levels, the basis for determining is the number of the period of the element.
- We divide the levels into sub levels.
Below you can see an example of how to correctly compose the electronic formulas of chemical elements.
- first we fill in the s-sublevel, and then the p-, d- b f-sublevels;
- according to the Klechkovsky rule, electrons fill the orbitals in the order of increasing energy of these orbitals;
- according to Hund's rule, electrons within one sublevel occupy free orbitals one at a time, and then form pairs;
- according to the Pauli principle, there are no more than 2 electrons in one orbital.
The electronic formula of a chemical element shows how many electronic layers and how many electrons are contained in an atom and how they are distributed over the layers.
To draw up the electronic formula of a chemical element, you need to look at the periodic table and use the information obtained for this element. The ordinal number of an element in the periodic table corresponds to the number of electrons in an atom. The number of electronic layers corresponds to the number of the period, the number of electrons on the last electronic layer corresponds to the number of the group.
It must be remembered that the first layer has a maximum of 2 electrons 1s2, the second has a maximum of 8 (two s and six p: 2s2 2p6), and the third has a maximum of 18 (two s, six p, and ten d: 3s2 3p6 3d10).
For example, the electronic formula of carbon: С 1s2 2s2 2p2 (serial number 6, period number 2, group number 4).
Electronic formula of sodium: Na 1s2 2s2 2p6 3s1 (ordinal number 11, period number 3, group number 1).
To check the correctness of the spelling of the electronic formula, you can look at the website www.alhimikov.net.
At first glance, drawing up an electronic formula for chemical elements may seem like a rather complicated task, but everything will become clear if you adhere to the following scheme:
- write the orbitals first
- insert numbers in front of the orbitals that indicate the number of the energy level. Do not forget the formula for determining the maximum number of electrons at the energy level: N = 2n2
How do you know the number of energy levels? Just look at the periodic table: this number is equal to the number of the period in which this element is located.
- above the orbital icon, write a number that indicates the number of electrons that are in this orbital.
For example, the electronic formula for scandium would look like this.
The task of drawing up the electronic formula of a chemical element is not the easiest one.
So, the algorithm for compiling electronic formulas of elements is as follows:
Here are the electronic formulas for some of the chemical elements:
You need to compose the electronic formulas of chemical elements in this way: you need to look at the number of the element in the periodic table, thus finding out how many electrons it has. Then you need to find out the number of levels, which is equal to the period. Then the sublevels are written and filled in:
First of all, you need to determine the number of atoms according to the periodic table.
To draw up an electronic formula, you will need the periodic system of Mendeleev. Find your chemical element there and see the period - it will be equal to the number of energy levels. The group number will correspond numerically to the number of electrons in the last level. The number of an element will be quantitatively equal to the number of its electrons. Also, you clearly need to know that at the first level there are a maximum of 2 electrons, at the second - 8, at the third - 18.
These are the highlights. In addition, on the Internet (including our website) you can find information with a ready-made electronic formula for each element, so you can check yourself.
Compilation of electronic formulas of chemical elements is a very difficult process, you cannot do without special tables, and a whole bunch of formulas must be used. In short, to compile, you need to go through these stages:
It is necessary to draw up an orbital diagram, in which there will be the concept of the difference between electrons from each other. Orbitals and electrons are highlighted in the diagram.
Electrons are filled in levels, from bottom to top, and have several sublevels.
So, first we find out the total number of electrons of a given atom.
We fill in the formula according to a certain scheme and write it down - this will be the electronic formula.
For example, for Nitrogen, this formula looks like this, first we deal with electrons:
And we write down the formula:
To understand the principle of drawing up the electronic formula of a chemical element, first you need to determine the total number of electrons in the atom by the number in the periodic table. After that, you need to determine the number of energy levels, taking as a basis the number of the period in which the element is located.
After that, the levels are divided into sublevels, which are filled with electrons, based on the Principle of Least Energy.
You can check the correctness of your reasoning by looking, for example, here.
Having compiled the electronic formula of a chemical element, you can find out how many electrons and electronic layers are in a particular atom, as well as the order of their distribution over the layers.
To begin with, we determine the ordinal number of the element according to the periodic table, it corresponds to the number of electrons. The number of electron layers indicates the number of the period, and the number of electrons on the last layer of the atom corresponds to the group number.
>> Chemistry: Electronic configurations of atoms of chemical elements
The Swiss physicist W. Pauli in 1925 established that in an atom in one orbital there can be no more than two electrons having opposite (antiparallel) spins (translated from English as "spindle"), that is, possessing such properties that can be conventionally 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 two, then these are paired electrons, that is, electrons with opposite spins.
Figure 5 shows a diagram of the subdivision of energy levels into sublevels.
The s-orbital, as you already know, is spherical. The electron of the hydrogen atom (s = 1) is located on this orbital and is unpaired. Therefore, its electronic formula or electronic configuration will be written as follows: 1s 1. In electronic formulas, the number of the energy level is indicated by the number in front of the letter (1 ...), the Latin letter denotes the sublevel (type of orbital), and the number written to the upper right of the letter (as an exponent) shows the number of electrons on the sublevel.
For a helium atom He, which has two paired electrons in one s-orbital, this formula is: 1s 2.
The electron shell of the helium atom is complete and very stable. Helium is a noble gas.
At the second energy level (n = 2), there are four orbitals: one s and three p. The electrons of the s-orbitals of the second level (2s-orbitals) have a higher energy, since they are at a greater distance from the nucleus than the electrons of the 1s-orbital (n = 2).
In general, for each value of n, there is one s-orbital, but with a corresponding store of electron energy on it and, therefore, with a corresponding diameter that grows as the value of n increases.
r-Orbital has the shape of a dumbbell or volumetric figure eight. All three p-orbitals are located in the atom mutually perpendicular along the spatial coordinates drawn through the nucleus of the atom. It should be emphasized once again that each energy level (electron layer), starting from n = 2, has three p-orbitals. With an increase in the value of n, electrons animate p-orbitals located at large distances from the nucleus and directed along the x, y, r axes.
For elements of the second period (n = 2), first one p-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 torn off - Be 0 is oxidized to the Be 2+ cation.
The fifth electron of the boron atom is occupied by a 2p orbital: 1s 2 2s 2 2p 1. Next, the C, N, O, E atoms are filled with 2p-orbitals, which ends in the noble gas of neon: 1s 2 2s 2 2p 6.
For the elements of the third period, the Sv and 3p orbitals are filled, respectively. In this case, five d-orbitals of the third level remain free:
11 Na 1s 2 2s 2 Sv1; 17S11v22822r63r5; 18Ag P ^ Ep ^ Zp6.
Sometimes in the 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 the atoms of chemical elements, in contrast to the above full electronic formulas.
In 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 enter the previous 3d and 4d orbitals, respectively (for elements of side subgroups): 23 V 2, 8, 11, 2; 26 Tr 2, 8, 14, 2; 40 Zr 2, 8, 18, 10, 2; 43 Tg 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 unfinished seventh - the electronic levels and sublevels are filled with electrons, as a rule, as follows: the first two electrons will enter the outer B-sublevel: 56 Ва 2, 8, 18, 18, 8, 2; 87Gg 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 enter the third outside energy level on 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 full filling with ten electrons, this level-equal will again be filled with the outer p-sublevel:
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 - the so-called graphic electronic formulas are written. For this notation, the following notation is used: each quantum cell is designated by a cell that corresponds to one orbital; each electron is indicated by an arrow corresponding to the direction of the spin. When writing a graphic electronic formula, two rules should be remembered: Pauli's principle, according to which there can be no more than two electrons in a cell (orbital), but with antiparallel spins, and F. Hund's rule, according to which electrons occupy free cells (orbitals), are located in they first one at a time and have the same spin value, and only then pair, but the spins, according to the Pauli principle, will already be oppositely directed.
In conclusion, we will once again consider the display of the electronic configurations of atoms of elements according to the periods of the DI Mendeleev system. Diagrams of the electronic structure of atoms show the distribution of electrons over the electron layers (energy levels). 
In a helium atom, the first electron layer is complete - there are 2 electrons in it.
Hydrogen and helium are s-elements, the s-orbital of these atoms is filled with electrons.
Elements of the second period
For all elements of the second period, the first electron layer is filled and 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 Pauli and Hund rules (Table 2).
In the neon atom, the second electron layer is complete - it contains 8 electrons.
Table 2 The structure of the electron shells of atoms of the elements of the second period
The end of the table. 2 
Li, Be - B-elements.
B, C, N, O, F, Ne - p-elements, these atoms are filled with electrons of the p-orbital.
Elements of the third period
For atoms of elements of the third period, the first and second electron layers are completed; therefore, the third electronic layer is filled, in which electrons can occupy the Зs-, 3p- and Зd-sublevels (Table 3).
Table 3 The structure of the electron shells of atoms of the elements of the third period
The 3s-electron orbital is being completed at the magnesium atom. Na and Mg-s-elements. 
The argon atom has 8 electrons on the outer layer (third electron layer). As the outer layer, it is complete, but in total in the third electron layer, as you already know, there may be 18 electrons, which means that the elements of the third period have Zd-orbitals unfilled.
All elements from Al to Ar are p-elements. s- and p-elements form the main subgroups in the Periodic Table.
For potassium and calcium atoms, a fourth electron layer appears, 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 in 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 3-elements. They belong to side subgroups, their pre-external electronic layer is filled, and they are referred to as transition elements.
Pay attention to the structure of the electron shells of chromium and copper atoms. In them there is a "dip" of one electron from the 4th to the 3rd sublevel, which is explained by the higher energy stability of the resulting electronic configurations Зd 5 and Зd 10: 
In the zinc atom, the third electronic layer is complete - all the sublevels 3s, Zp and Zd are filled in it, with a total of 18 electrons on them.
In the elements following zinc, the fourth electronic layer, the 4p-sublevel, continues to be filled: Elements from Ga to Kr are p-elements. 
At the krypton atom, the outer layer (fourth) is complete, it has 8 electrons. But in total in the fourth electron layer, as you know, there can be 32 electrons; for the krypton atom, the 4d and 4f sublevels are still empty.
For the elements of the fifth period, the sublevels are filled in the following order: 5s-> 4d -> 5p. And there are also exceptions associated with the "dip" 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 electron layer are filled, respectively.
The 4f-elements are called lanthanides.
5f-Elements are called actinides.
The order of filling the electronic sublevels in the atoms of the elements of the sixth period: 55 Сs and 56 Ва - 6s-elements;
57 Lа ... 6s 2 5d 1 - 5d-element; 58 Ce - 71 Lu - 4f-elements; 72 Hf - 80 Hg - 5d elements; 81 Тl- 86 Rn - 6p-elements. But even here there are elements in which the order of filling of electron orbitals is "violated", which, for example, is associated with a higher energy stability of half and fully 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; filled with electrons in the sublevel of the outer level of the atom; 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 pre-outer level of the atom is filled with electrons; d-elements include elements of secondary subgroups of groups I-VIII, that is, elements of inserted decades of large periods located between s- and p-elements. They are also called transition elements;
4) f-elements, filled with electrons f-sublevel of the third outside the level of the atom; these include lanthanides and actinides.
1. What would happen if the Pauli principle was not observed?
2. What would happen if Hund's rule was not followed?
3. Make diagrams of the electronic structure, electronic formulas and graphic electronic formulas of the atoms of the following chemical elements: Ca, Fe, Zr, Sn, Nb, Hf, Pa.
4. Write the electronic formula for element 110 using the symbol for the corresponding noble gas.
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