Describe the electronic configurations of atoms 19. Electronic configurations of atoms of the elements of the Periodic system. Distribution of electrons using the periodic system of D. I. Mendeleev

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 total 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 familiar with.

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 the 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 atom of iron
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).

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.

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 group VIIIA (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 general 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 that 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 group, which includes this element, 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. 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 ?

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).

Previously, we made up 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, as you remember, they have a 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

Algorithm for drawing up the electronic formulas of atoms (for example, the iodine atom)

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 COMPOSING 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.

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 of 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 have equal number of valence electrons no matter what orbitals these electrons are in. So, the elements 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, fall here in 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 have constructed 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?

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 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").
Please 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?

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.

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 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?

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