The electronic formula of sulfur is zero. Electronic configurations of atoms of chemical elements - Knowledge Hypermarket

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 for atoms of some elements, the sequence of filling orbitals with electrons is violated. Sometimes these violations are called "exceptions", but this is not so - there are no exceptions to the laws of Nature!

The first element with such a violation is chromium. Let us consider in more detail its electronic structure (Fig. 6.16 A). The chromium atom has 4 s-sublevel is not two, as one would expect, but only one electron. But for 3 d-sublevel five electrons, but this sublevel is filled after 4 s-sublevel (see Fig. 6.4). To understand why this happens, let's look at what electron clouds are 3 d sublevel of this atom.

Each of the 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 is spherical, or, as they say, spherically symmetrical. By the nature of the electron density distribution in different directions, it is similar to 1 s-EO. The energy of the sublevel whose electrons form such a cloud turns out to be lower than in the case of a less symmetrical cloud. In this case, the energy of orbitals 3 d-sublevel is equal to energy 4 s-orbitals. When the symmetry is broken, for example, when the sixth electron appears, the energy of the orbitals is 3 d-sublevel again becomes more than energy 4 s-orbitals. Therefore, the manganese atom again has a second electron for 4 s-AO.
Spherical symmetry has a common cloud of any sublevel filled with electrons both half and completely. 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 must look for the next violation in the atom, in the electron shell of which the ninth “comes” last d-electron. Indeed, the copper atom has 3 d-sublevel 10 electrons, and 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 during the interaction of the electron shells of atoms. For all atoms (except hydrogen), not all electrons take part in the formation of chemical bonds: for boron, three out of five electrons, for carbon, four out of six, and, for example, for barium, two out of fifty-six. These "active" electrons are called valence electrons.

Sometimes valence electrons are confused with external electrons, but they are not the same thing.

The electron clouds of 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 bonds 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 along with them, part of the electrons can also take part in the formation of a bond. pre-external(penultimate) layer, but only if they have an energy not much different from the energy of the outer electrons. Both those and other electrons of the atom are valence. (In lanthanides and actinides, even some "pre-external" electrons are valence)
The energy of valence electrons is much greater than the energy of other electrons of the atom, and valence electrons differ much less in energy from each other.
Outer 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 whose electronic configuration 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, they 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 the sublevels formed by these orbitals, that is, 4 s-, 4p-, 4d- and 4 f-EPU.
Outer electrons are always valence, therefore, 4 s-electrons of an iron atom are valence electrons. And if so, then 3 d-electrons with a 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. All of them are external, but only 4 are valence R-AO, since the energy of the remaining 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 - the fourth,
outer 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),
the outer electron layer is the fourth,
external electron 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).

Valence sublevels can be partially or completely filled with electrons, or they can remain free at all. With an increase in the charge of the nucleus, the energy values ​​of all sublevels decrease, but due to the interaction of electrons with each other, the energy of different sublevels decreases with different "speed". The energy of fully 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 titanium atom 3 d-EPU is only partially filled with electrons, and its energy is greater than the energy of 4 s-EPU, and 3 d-electrons are valence. At the arsenic atom 3 d-EPU is completely filled with electrons, and its energy is much less than energy 4 s-EPU, and therefore 3 d-electrons are not valence.
In these examples, we analyzed valence electronic configuration titanium and arsenic atoms.

The valence electronic configuration of an atom is depicted as valence electronic formula, or in the form energy diagram of valence sublevels.

VALENCE ELECTRONS, EXTERNAL ELECTRONS, VALENCE EPU, VALENCE AO, VALENCE ELECTRON CONFIGURATION OF THE ATOM, VALENCE ELECTRON FORMULA, VALENCE SUBLEVEL DIAGRAM.

1. On the energy diagrams you have compiled and in the full electronic formulas of the atoms Na, Mg, Al, Si, P, S, Cl, Ar, indicate the external and valence electrons. Make up the valence electronic formulas 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 are 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 at the atom of a) boron, b) fluorine, c) sodium?
5. How many orbitals with an unpaired electron does an atom have a) boron, b) fluorine, c) iron
6. How many free outer orbitals does a manganese atom have? 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 series of elements to this strip (from hydrogen to meitnerium).
8. In each cell, place the symbol of the element, its serial 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 electron shells their atoms.
With a natural side chemical elements you are already familiar. Now let's get acquainted with the principle of similarity of electron 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 main 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 their corresponding valence configurations can be called similar. For atoms of different elements with similar electronic configurations, we can write common 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, in the natural series of elements, atoms with similar valence electronic configurations regularly occur. This is the principle of similarity of electron shells.
Let us try to reveal the form of this regularity. To do this, we will use the natural series of elements you made.

NRE begins with hydrogen, whose valence electronic formula is 1 s 1 . In search of similar valence configurations, we cut the natural series of elements in front of elements with a common valence electronic formula ns 1 (that is, before lithium, before sodium, etc.). We have received so-called "periods" of elements. Let's add the resulting "periods" so that they become table rows (see Figure 6.20). As a result, only the atoms of the first two columns of the table will have such electronic configurations.

Let's try to achieve similarity of valence electronic configurations in other columns of the table. To do this, we cut out elements with numbers 58 - 71 and 90 -103 from the 6th and 7th periods (they have 4 f- and 5 f-sublevels) and place them under the table. The symbols of the remaining elements will be shifted horizontally as shown in the figure. After that, the atoms of the elements in the same column of the table will have similar valence configurations, which can be expressed in 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 explained by the same reasons as in the case of chromium and copper (see paragraph 6.6).

As you can see, using the NRE and applying the principle of similarity of electron shells, we managed to systematize the chemical elements. Such a system of chemical elements is called natural, as it is based solely on the laws of Nature. The table we received (Fig. 6.21) is one of the ways to graphically depict a 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 get acquainted in more detail with the structure of the long-period table of chemical elements.
The rows of this table, as you already know, are called "periods" of the elements. Periods are numbered with Arabic numerals from 1 to 7. There are only two elements in the first period. The second and third periods, containing eight elements each, are called short periods. The fourth and fifth periods, containing 18 elements each, are called long periods. The sixth and seventh periods, containing 32 elements each, are called extra long periods.
The columns of this table are called groups elements. Group numbers are indicated by Roman numerals with Latin letters A or B.
The 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)(names "alkali metals" and alkaline earth metals" refer to simple substances formed by the respective elements and should not be used as names of groups of elements); group VIA elements (O, S, Se, Te, Po) - chalcogens, elements of group VIIA (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 applies to simple substances)
The elements usually placed in the lower part of the table with serial numbers 58 - 71 (Ce - Lu) are called lanthanides("following lanthanum"), and elements with serial numbers 90 - 103 (Th - Lr) - actinides("following actinium"). There is a variant of the long-period table, in which the lanthanides and actinides are not cut out of the NRE, but remain in their places in extra-long periods. This table is sometimes called extra long period.
The long period table is divided into four block(or sections).
s-block includes elements of IA and IIA groups with common valence electronic formulas ns 1 and ns 2 (s-elements).
p-block includes elements from group IIIA to VIIIA 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 common 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).

Elements s- And p-blocks form A-groups, and elements d-block - B-group of a system of chemical elements. All f-elements are formally included in group IIIB.
The elements of the first period - hydrogen and helium - are s-elements and can be placed in IA and IIA groups. But helium is more often placed in group VIIIA as the element with which the period ends, which is fully consistent with its properties (helium, like all other simple substances formed by the elements of this group is a noble gas). Hydrogen is often placed in group VIIA, since its properties are much closer to halogens than to alkaline elements.
Each of the periods of the system begins with an element that has 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 the valence configuration of atoms ns 2 np 6 (except for the first period). This makes it easy to identify groups of sublevels in the energy diagram that are filled with electrons at the atoms of each of the periods (Fig. 6.22). Do this work with all the sublevels shown in the copy you made of Figure 6.4. The sublevels highlighted in Figure 6.22 (except for fully 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 are fully consistent with the sequence of filling s-, p-, d- or f- sublevels of electrons. This feature of the system of elements allows, knowing the period and group, which includes a given element, to immediately write down its valence electronic formula.

LONG-PERIOD TABLE OF CHEMICAL ELEMENTS, BLOCKS, PERIODS, GROUPS, ALKALINE ELEMENTS, ALKALINE EARTH ELEMENTS, CHALCOGENES, HALOGENS, NOBLE GAS ELEMENTS, LANTHANOIDES, ACTINOIDES.
Write down the general valence electronic formulas of the atoms of the 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 in the direction of 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 atoms 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 an atom belongs 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, make valence, complete and abbreviated electronic formulas and draw energy diagrams of valence sublevels. 9. The 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 different purposes, we need to know either the full or valence configuration of an atom. Each of these electronic configurations can be represented both by a formula and by an energy diagram. That is, complete electronic configuration of an atom expressed the full electronic formula of the atom, or full energy diagram of an atom. In its turn, valence electron configuration of an atom expressed valence(or, as it is often called, " short ") the electronic formula of the atom, or diagram of valence sublevels of an atom(Fig. 6.23).

Previously, we made electronic formulas of atoms using the ordinal numbers of the elements. At the same time, we determined the sequence of filling 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 and so on. And only by writing down the full electronic formula, we could also write down the valence formula.
It is more convenient to write the valence electronic formula of the atom, which is most often used, based on the position of the element in the system of chemical elements, according to the period-group coordinates.
Let's consider in detail how this is done for elements s-, p- And d-blocks.
For elements s-block valence electronic formula of an atom consists of three characters. In general, it can be written like this:

In the first place (in the place of a large cell) is the period number (equal to the main quantum number of these s-electrons), and on the third (in the superscript) - the number of the group (equal to the number of valence electrons). Taking as an example a magnesium atom (3rd period, group IIA), we get:

For elements p-block valence electronic formula of an atom consists of six characters:

Here, in place 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, VIA group) we get:

2s 2 2p 4 .

Valence electronic formula of most elements d block can be written like this:

As in 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. An example is 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-sublevel has only one electron, and the general valence electronic formula ns 1 (n–1)d 5 . Therefore, the valence electronic formula, for example, of molybdenum (5th period) is 5 s 1 4d 5 .
It is also easy to make a valence electronic formula of any element of the IB group, for example, gold (6th period)>–>6 s 1 5d 10 , but in this case you need to remember 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 10 . Therefore, the valence electronic formula, for example, of a zinc atom is 4 s 2 3d 10 .
General rules the valence electronic formulas of the elements of the first triad (Fe, Co and Ni) also obey. Iron, an element of group VIIIB, has a valence electronic formula of 4 s 2 3d 6. The cobalt atom has one d-electron more (4 s 2 3d 7), while the nickel atom has 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 in them, due to the tendency to highly symmetric electron shells, the filling of valence sublevels with electrons has some additional features.
Knowing the valence electronic formula, one can also write down the complete electronic formula of the atom (see below).
Often, instead of cumbersome full electronic formulas, they write down abbreviated electronic formulas atoms. To compile them in the electronic formula, all the electrons of the atom except 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 forming the noble gas) is replaced by 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 compiling electronic formulas of atoms (on the example of an iodine atom)

Notes
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 the IB group.

COMPLETE ELECTRONIC FORMULA, VALENCE ELECTRONIC FORMULA, abbreviated ELECTRONIC FORMULA, ALGORITHM FOR COMPOSING ELECTRONIC FORMULA OF ATOMS.
1. Compose the valence electronic formula of the atom of the element a) the second period of the third A group, b) the third period of the second A group, c) fourth period fourth A group.
2. Make abbreviated electronic formulas of magnesium, phosphorus, potassium, iron, bromine and argon atoms.

Over the 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 that graphically reflect this system. Of these, in addition to the long-period table, the so-called short-period table of elements of D. I. Mendeleev is most widely used. A short-period table is obtained from a long-period one, if the 4th, 5th, 6th and 7th periods are cut in front of the elements of the IB group, moved apart and the resulting rows are added in the same way as we added the periods before. The result is shown in figure 6.24.

The lanthanides and actinides are also placed under the main table here.

IN groups this table contains elements whose atoms have the same 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 similarity of electron shells, fall here into the same seventh group. The need to distinguish between such elements makes it necessary to single out in groups subgroups: main- analogues of A-groups of the long-period table and side effects are analogues of B-groups. In Figure 34, the symbols of the elements of the main subgroups are shifted to the left, and the symbols of the elements of the secondary subgroups are shifted to the right.
True, such an 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 laws of the electronic structure of atoms, the similarity and patterns 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 system of elements, and much more. The short period table is less convenient in this respect.

SHORT-PERIOD TABLE, MAIN SUB-GROUPS, SECONDARY SUB-GROUPS.
1. Convert the long-period table you built from the natural series of elements into a short-period table. Carry out the reverse transformation.
2. Is it possible to make 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 a radius r oo. 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 an 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. Consequently, the orbital radius of an atom also depends on these same characteristics.
Consider the electron shells of hydrogen and helium atoms. Both in the hydrogen atom and in the helium atom, electrons are located on 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 that of 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. Therefore, the radius of a helium atom must be much smaller than the radius of a hydrogen atom. This is true: r 0 (He) / r 0 (H) \u003d 0.291 E / 0.529 E 0.55.
The lithium atom has an outer electron at 2 s-AO, that is, 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 placed in the same second electron layer, and the charge of the nucleus of these atoms increases with increasing serial number. Electrons are more strongly attracted to the nucleus, and, naturally, the radii of atoms decrease. We could repeat these arguments for the atoms of the elements of other periods, but with one clarification: the orbital radius monotonically decreases 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 a system of elements is as follows: with an increase in the serial number in a period, the orbital radii of atoms decrease, and in a group they increase. The largest atom is a cesium atom, and the smallest is a helium atom, but of the atoms of the 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, standing in the natural series 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, consequently, the nuclear charge 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 be attracted in the absence of lanthanides (this effect is often called "lanthanide contraction").
Please note that when passing from atoms of elements of group VIIIA to atoms of elements of group IA, 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 the dependence of the orbital radius of the atom on the element's serial number 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 give an electron in an atom additional energy (you will learn how to do this from a physics course), then the electron can go to another AO, that is, the atom will end up in excited state. This state is unstable, and the electron will almost immediately return to its original state, and excess energy will be released. But if the energy imparted to the electron is large enough, the electron can completely break away from the atom, while the atom ionized, that is, it turns into a positively charged ion ( cation). The energy needed to do this is called ionization energy of an atom(E And).

It is quite difficult to tear off 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).

Molar ionization energy shows what is the smallest energy 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 the element in the system of elements, that is, how does it change in the group and period?
In physical terms, the ionization energy is equal to the work that must be expended to overcome the force of attraction of an electron to an atom when moving an electron from an atom to an infinite distance from it.

Where q is the charge of an electron, Q is the charge of the cation remaining after the removal of an electron, and r o is the orbital radius of the atom.

AND q, And Q are constant values, and it can be concluded that, the work of detaching an electron A, and with it the ionization energy E and, are inversely proportional to the orbital radius of the atom.
After analyzing the values ​​of the orbital radii of atoms of various elements and the corresponding values ​​of the ionization energy given in Appendixes 5 and 6, you can see that the relationship between these values ​​is close to proportional, but somewhat different from it. The reason that our conclusion does not agree well with the experimental data is that we used a very rough 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 atoms decreases in a period with an increase in the serial number, the ionization energy increases. In a group, as the atomic number increases, the orbital radius of the atoms, as a rule, increases, and the ionization energy decreases. The highest molar ionization energy is in the smallest atoms, helium atoms (2372 kJ/mol), and of 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 ionization energy can be schematically shown as follows:

In chemistry, it is important that the ionization energy characterizes the propensity of an atom to donate "its" electrons: the greater the ionization energy, the less inclined the atom is to donate electrons, and vice versa.

Excited state, ionization, cation, ionization energy, molar ionization energy, change in ionization energy in a system of elements.
1. Using the data given in Appendix 6, determine how much energy you need to spend to tear off one electron 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 does this ratio differ from the ratio of the molar ionization energies of the same atoms?
3. According to the data given in Appendix 6, plot 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 for the previous paragraph. See if this graph matches the choice 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 ionization energy, the corresponding molar quantity is usually used - molar electron affinity energy().

The molar electron affinity energy shows what is the energy released when one mole of electrons is added to one mole of neutral atoms (one electron to each atom). Like the molar ionization energy, this quantity 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 the atom, the electron, it would seem, should be repelled by the same negatively charged electrons that form the electron shell. Actually this is not true. Remember if you have ever dealt with atomic chlorine. Of course not. After all, it exists only at very high temperatures. Even more stable molecular chlorine is practically not found in nature - if necessary, it has to be obtained using chemical reactions. And you have to deal with sodium chloride (common salt) all the time. After all, table salt is consumed by a person with food every day. And it is quite common in nature. But after all, table salt contains chloride ions, that is, chlorine atoms that have attached one "extra" electron each. 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.
Due to the fact that several factors influence the value of the electron affinity energy, the nature of the change in this value in a 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 in general terms similar to a change in the ionization energy, that is, the energy of electron affinity in the group decreases, and in the period it increases. It is maximum at the atoms of fluorine (328 kJ/mol) and chlorine (349 kJ/mol). The nature of the change in the electron affinity energy in the system of elements resembles the nature of the change in the ionization energy, that is, the direction of the increase in the electron affinity energy 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 electron affinity on the serial number for atoms of elements with Z from 1 to 40 using app 7.
3.What physical meaning have negative electron affinity energies?
4. Why, of all the atoms of the elements of the 2nd period, only beryllium, nitrogen and neon have negative values ​​of the molar energy of electron affinity?

You already know that the propensity of an atom to donate its own and accept foreign electrons depends on its energy characteristics (ionization energy and electron affinity energy). What atoms are more inclined to donate their electrons, and which ones are more inclined to accept strangers?
To answer this question, let us summarize in Table 15 everything that we know about the change in these inclinations in the system of elements.

Table 15