Na electrons. Electronic atom configurations. Electronic structure of an oxygen atom

6.6. Features of the electronic structure of chromium, copper atoms and some other elements

If you carefully looked at Appendix 4, then, probably, we noticed that at the atoms of some elements, the sequence of filling electrons of orbitals is broken. Sometimes these violations are called "exceptions", but it is not so - no exceptions from the laws of nature does not happen!

The first element with such a disorder is chrome. Consider more of its electronic structure (Fig. 6.16 but). At the chromium atom on 4 s.-Provers are not two, as it should be expected, but only one electron. But on 3. d.-provine five electrons, but this sublevel is filled after 4 s.-Production (see Fig. 6.4). To understand why this happens, let's see what the electronic clouds 3 are represented d.-Production of this atom.

Each of five 3 d.- Poland in this case is formed by one electron. As you already know from § 4 of this chapter, the general electronic cloud of such five electrons has a spherical shape, or, as they say, spherically symmetrically. By the nature of the distribution of electron density in different directions, it looks like 1 s.-Eho. The energy of a suite, 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 3 d.-Production equal to energy 4 s.-Ibed. In case of symmetry violation, for example, when the sixth electron appears, the energy of the orbital d.-Production again becomes more than energy 4 s.-Ibed. Therefore, at the manganese atom, the second electron appears again on 4 s.-Ao.
Sphoric symmetry has a common cloud of any subproduction filled with electrons both half and completely. A decrease in energy in these cases is common and does not depend on whether by half or completely filled with electrons of any kind. And if so, the next violation we should look for at the atom, in the electronic shell of which the last "comes" the ninth d.-electron. And indeed, at the copper atom on 3 d.-provine 10 electrons, and on 4 s.-Provers only one (Fig. 6.16 b.).
Reducing the energy of the orbitals is fully or half the filled sublevel is the cause of a number of important chemical phenomena, with some of which you will also get acquainted.

In the chemistry, the properties of isolated atoms, as a rule, are not studied, since almost all atoms, entering the composition of various substances, form chemical bonds. Chemical bonds are formed in the interaction of electronic shells of atoms. All atoms (except hydrogen) are not at the formation of chemical ties, not all electrons take part: Bora is three electrons out of five, carbon is four out of six, and, for example, barium is two out of fifty six. These "active" electrons are called valence electrons.

Sometimes valence electrons are confused with externalelectrons, and this is not the same thing.

Electronic clouds of external electrons have a maximum radius (and the maximum value of the main quantum number).

It is the external electrons that take part in the formation of communication in the first place, if only because, with rapprochement of atoms, electronic clouds formed by these electrons come into contact above all. But together with them, part of electrons can also participate in education antichenous(penultimate) layer, but only if they have an energy that does not differ from the energy of external electrons. And those and other electrons of the atom are valence. (Lantanoids and actinides are even some "bridal" electrons)
The energy of valence electrons is much larger than the energy of other electrons of the atom, and each other valence electrons by energy differ significantly less.
External electrons are always valence only if the atom can form chemical bonds in general. Thus, both electrons of the helium atom are external, but it is impossible to name their valence, since the helium atom does not form any chemical bonds.
Valented electrons occupy valental orbitalswhich in turn form valenny Slovennel.

As an example, consider an iron atom, the electronic configuration of which is shown in Fig. 6.17. From the electrons of the iron atom, the maximum main quantum number ( n.\u003d 4) have only two 4 s.-Electron. Consequently, they are the external electrons of this atom. Outer orbitals of iron atom - all orbitals with n. \u003d 4, and external sublocks - all the sublevels formed by these orbital, i.e. 4 s.-, 4p.-, 4d.- and 4. f.-Pep.
External electrons - always valence, therefore 4 s.-Electrons atom of iron - valence electrons. And if so, then 3 d.-Electrons having a little big energy will also be valence. At the external level of the iron atom except completed 4 s.-Ao there are still free 4 p.-, 4d.- and 4. f.-Ao. All of them external, but valence among them only 4 r-Ao, since the energy of the remaining orbitals is much larger, and the appearance of electrons on these orbitals for an iron atom is not profitable.

So, at the iron atom
External electronic level - fourth,
External Sloves - 4 s.-, 4p.-, 4d.- and 4. f.-EPU,
External orbitals - 4 s.-, 4p.-, 4d.- and 4. f.-Ao
External electrons - two 4 s.-Electron (4. s. 2),
External electronic layer - fourth,
External electronic cloud - 4 s.-EEO
Valentines Sloves - 4 s.-, 4p.-, and 3 d.-EPU,
Valental 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 ties can be filled with electrons partially or completely, and can and remain free. With an increase in the charge of the kernel, the energy values \u200b\u200bof all supremes are reduced, but due to the interaction of electrons among themselves, the energy of different sublevels decreases with different "speed". Energy completely filled d.- I. f.-proving is reduced so much that they cease to be valence.

As an example, consider the titanium and arsenic atoms (Fig. 6.18).

In the case of a titanium atom 3 d.-Epople is filled with electrons only partially, and its energy is greater than energy 4 s.-EPU, and 3 d.- Electrons are valence. Arsenic atom has 3 d.-EPU completely filled with electrons, and its energy is significantly less than energy 4 s.-EPU, and, therefore, 3 d.-Electrons are not valence.
In the examples, we analyzed valenny electronic configurationtitanium and arsenic atoms.

The valence electronic atom configuration is depicted as valence electronic formulaor in the form of energy chart of valence pins.

Valence electrons, external electrons, valence eppa, valence AO, valence electronic configuration of an atom, valence electronic formula, valence diagram of iron.

1. The energy diagrams compiled by you and in full electronic formulas of Na, Mg, Al, Si, P, S, S, Cl, AR atoms include external and valence electrons. Make the valence electronic formulas of these atoms. On energy diagrams, highlight parts corresponding to the energy diagrams of valence pins.
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 much valenny sublevels in the electronic shell of the 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 with atom a) boron, b) fluorine, c) sodium?
5. How much orbitals with an unpaired electron at an atom a) boron, b) fluorine, c) iron
6. How much free external orbital at the manganese atom? And how many free valenny?
7. For the next session, prepare a strip of paper 20 mm wide, divide it on cells (20? 20 mm), and apply a natural row of elements to this strip (from hydrogen to a pointener).
8. In each cell, place the symbol of the element, its sequence number and the valence electronic formula, as shown in Fig. 6.19 (Use Appendix 4).

The systematization of chemical elements is based on a natural range of elements. and principle of similarity of electronic shellstheir atoms.
With a natural number of chemical elements you are already familiar. Now let's get acquainted with the principle of similarity of electronic shells.
Considering the valence electronic formulas of atoms in ERE, it is easy to detect that in some atoms they differ only with the values \u200b\u200bof the main quantum number. For example, 1. s. 1 at hydrogen, 2 s. 1 in lithium, 3 s. 1 at sodium, etc. or 2 s. 2 2p. 5 at Fluoro, 3 s. 2 3p. 5 Chlorine, 4 s. 2 4p. 5 in bromine, etc. This means that the outer areas of the clouds of valence electrons of such atoms are very similar and differ only in size (and, of course, electron density). And if so, the electronic clouds of such atoms and the corresponding valet configurations can be called similar. For atoms of different elements with similar electronic configurations, we can record general valence electronic formulas: nS. 1 in the first case and nS. 2 nP. 5 in the second. Moving along the natural row of elements, you can find other groups of atoms with similar valet configurations.
In this way, in the natural row of elements, atoms with similar valence electronic configurations are regularly found.. This is the principle of similarity of electronic shells.
Let's try to identify the view of this regularity. To do this, we use the natural number of items made by you.

Ere begins with hydrogen, whose valence electronic formula 1 s. one . In search of such valet configurations we will cut the natural range of elements in front of elements with a total valence electronic formula nS. 1 (i.e., before lithium, before sodium, etc.). We got the so-called "periods" of the elements. Moving the resulting "periods" so that they become strings of the table (see Fig. 6.20). As a result, such electronic configurations will be only at the atoms of the first two columns of the table.

Let's try to achieve the similarity of valence electronic configurations and in other columns of the table. To do this, cut out from the 6th and 7th periods of elements with numbers 58 - 71 and 90-103 (they are filling 4 f.- and 5. f.-proving) and put them under the table. The characters of the remaining elements slide horizontally as shown in the figure. After that, atoms of elements facing one table of the table, such valence configurations will be obtained, which can be expressed by common valence electronic formulas: nS. 1 , nS. 2 , nS. 2 (n.–1)d. 1 , nS. 2 (n.–1)d. 2 and so on to nS. 2 nP. 6. All deviations from the total 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 ERE and applying the principle of similarity of electronic shells, we managed to systematize chemical elements. Such a system of chemical elements is called naturalSince it is based solely on the laws of nature. The table we obtained (Fig. 6.21) is one of the ways of graphic image of a natural system of elements and is called long-range table of chemical elements.

The principle of the similarity of electronic shells, the natural system of chemical elements ("periodic" system), a table of chemical elements.

We will get acquainted with the structure of the long-range table of chemical elements.
Lines of this table, as you already know, are called "periods" of elements. Periods are numbered by Arabic numbers from 1 to 7. In the first period there are only two elements. The second and third periods containing eight elements are called shortperiods. The fourth and fifth periods containing 18 elements are called longperiods. The sixth and seventh periods containing 32 elements are called overlongs Periods.
Columns of this table are called groupselements. Group numbers are indicated by Roman numbers with Latin letters A or V.
Elements of some groups have their general (group) names: elements of the IA group (Li, Na, K, RB, CS, FR) - alkaline elements(or elements of alkali metal); IIA elements of the group (CA, SR, BA and RA) - alkaline earth elements(or elements of alkaline earth metals) (The name "Alkali metals" and alkaline earth metals "belong to the simple substances formed by the corresponding elements and should not be used as the names of the groups of elements); elements of the group VIA (O, S, SE, TE, PO) - hallcohele, elements of the group VIIA (F, CL, BR, I, AT) - halogens, elements of the group VIIIA (HE, NE, AR, KR, XE, RN) - elements of noble gases. (The traditional name "noble gases" also belongs to simple substances)
Usually submitted to the bottom of the table items with sequence numbers 58 - 71 (CE - LU) are called lantanoids ("Next Lantane"), and elements with sequence numbers 90 - 103 (Th - LR) - aktinoids ("Next Activity"). There is a variant of a long-range table, in which lanthanoids and actinoids are not cut from Ere, but remain in their places in super long periods. Such a table is sometimes called supersensnoiodine.
The long-range table is divided into four blok(or sections).
s-block Includes IA elements and IIA groups with shared valence electronic formulas nS. 1 I. nS. 2 (s-elements).
r-block Includes elements with IIIa viiia group with common valence electronic formulas from nS. 2 nP. 1 BE nS. 2 nP. 6 (p-Elements).
d-block Includes elements with IIIB via IIB Group with common valence electronic formulas from nS. 2 (n.–1)d. 1 BE nS. 2 (n.–1)d. 10 (d-elements).
f-block Includes lanthanoids and actinoids ( f-elements).

Elements s.- I. p.-blocks form a-groups, and elements d. -block - B-groups of system of chemical elements. Everything f.- elements formally included in the IIIB group.
Elements of the first period - hydrogen and helium - are s.- elements and can be placed in IA and IIA groups. But helium is more likely placed in the group as an element that the period ends, which fully corresponds to its properties (helium, as well as all other substances formed by the elements of this group - noble gas). Hydrogen is often placed in the VIIA group, since in its properties it is essentially closer to halogens than alkaline elements.
Each of the period periods begins with an element having a valence configuration of atoms nS. 1, since it is precisely from these atoms that the formation of the next electronic layer begins, and ends with an element with valence configuration of atoms nS. 2 nP. 6 (except for the first period). This makes it easy to distinguish on the energy diagram of a group of sublevels filled with electrons at the atoms of each period (Fig. 6.22). Do this work with all the supremes depicted on the copies of Figure 6.4. Highlighted in Figure 6.22 of the subframe (except fully filled d.- I. f.-proving) are valence for atoms of all elements of this period.
Appearance in periods s.-, p.-, d.- or f.-Elements fully corresponds to the fill sequence s.-, p.-, d.- or f.-provine electrons. This feature of the system of elements allows, knowing the period and the group in which this element includes, immediately record its valence electronic formula.

Long-range table of chemical elements, blocks, periods, groups, alkaline elements, alkaline earth elements, chalcogen, halogens, elements of noble gases, Lantanoids, actinoids.
Write down the general valence electronic formulas of elements a) IVA and IVB groups, b) IIIA and VIIB groups?
2. What is common between the electronic configurations of the atoms of elements A and in groups? What do they differ?
3. How many groups of elements are included in a) s.-Block, b) r-block, c) d.-block?
4. Conduct Figure 30 in the direction of increasing the energy of the supreme and highlight the groups of sublevels filled with electrons in the 4th, 5th and 6th periods.
5. Transfer valence lint atoms a) calcium, b) phosphorus, c) titanium, d) chlorine, d) sodium. 6. Formulate how the S-, P- and D-elements differ from each other.
7. Equally, why the atom belongs to any element is determined by the number of protons in the nucleus, and not the mass of this atom.
8. For lithium, aluminum atoms, strontium, selenium, iron and lead, make up valence, complete and abbreviated electronic formulas and depict energy diagrams of valence pins. 9. What elements of the 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 complete or valence atom configuration. Each of these electronic configurations can be depicted as a formula and energy diagram. I.e, full electronic atom configurationexpress full electronic formula, or full energy diagram atom. In its turn, valence Electronic Atom Configurationexpress valentine(or, as it is often called, " brief) electronic formula atom, or diagram of valence sublayer atom(Fig. 6.23).

Previously, we accounted for electronic formulas atoms using the sequence numbers of the elements. At the same time, we determined the sequence of filling with sublayer electrons by energy diagram: 1 s., 2s., 2p., 3s., 3p., 4s., 3d., 4p., 5s., 4d., 5p., 6s., 4f., 5d., 6p., 7s. and so on. And only writing the complete electronic formula, we could record the valence formula.
The valence electronic formula of the atom, which is most often used, more convenient to record, based on the position of the element in the system of chemical elements, along the coordinates of the period - group.
Consider in detail how it is done for items. s.-, p.- I. d.-blocks.
For elements s.-block The valence electronic formula of the atom consists of three characters. In general, it can be written as:

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

For elements p.-block The valence electronic formula of the atom consists of six characters:

Here at the site of large cells is also raised by the number of the period (equal to the main quantum number of these s.- I. p.-Electrons), and the number of the group (equal to the number of valence electrons) turns out to be equal to the sum of the upper indices. For an 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 the previous cases, the number of the period is installed instead of the first cell (equal to the main quantum number of these s.-Electrons). The number in the second cell turns out to be less, since one is less than the main quantum number of these d.-Electrons. The group number here is also equal to the amount of indexes. Example - Valence electronic Titan formula (4th period, IVB Group): 4 s. 2 3d. 2 .

The group number is equal to the amount of indexes and for the VIB elements of the group, but they are like you remember on the valence s.-Provers only one electron, and general valence electronic formula nS. 1 (n.–1)d. five . Therefore, the valence electronic formula, for example, molybdenum (5th period) - 5 s. 1 4d. 5 .
It is also easy to make a valence electronic formula of any element IB group, for example, gold (6th period)\u003e -\u003e 6 s. 1 5d. 10, but in this case you need to remember that d.- electrons at the atoms of the elements of this group still remain valence, and some of them can participate in the formation of chemical bonds.
General valence electronic formula of atoms of elements of group IIB - nS. 2 (n. – 1)d. ten . Therefore, a valence electronic formula, for example, zinc atom - 4 s. 2 3d. 10 .
The overall rules are subject to valence electronic formulas of the elements of the first triad (Fe, CO and Ni). Iron, element VIIIB group, valence electronic formula 4 s. 2 3d. 6. At the cobalt atom - one d.-Electron more (4 s. 2 3d. 7), and at the nickel atom - for two (4 s. 2 3d. 8).
Using only by these rules for writing valence electronic formulas, it is impossible to compile electronic formulas of atoms of some d.-Elements (NB, RU, RH, PD, IR, PT), as they have due to the desire for highly-sensitive electronic shells filling with electron blasting sublevels have some additional features.
Knowing the valence electronic formula, and the complete electronic formula of the atom can be recorded (see below).
Often, instead of bulky complete electronic formulas write down abbreviated electronic formulasatoms. To compile them, in the electron formula, all electrons of an atom other than valence are placed in the square brackets and part of the electronic formula corresponding to the electronic formula of the atom of the last element of the preceding 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 the preparation of electronic formulas of atoms (on the example of an iodine atom)

Notes
1. For the 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 valented at the atoms of elements of the IB group.

Full electronic formula, valence electronic formula, abbreviated electronic formula, algorithm for compiling electronic formulas atoms.
1. Suggest the valence electronic formula of an element a) of the second period of the third A group, b) of the third period of the second A group, c) of the fourth period of the fourth A group.
2. Suggest the abbreviated electronic formulas of magnesium atoms, phosphorus, potassium, iron, bromine and argon.

For 100 years since the opening of the natural system of elements, several hundreds of a wide variety of tables that graphically reflect this system were proposed. Of these, except for the long-range table, the so-called short-range table of elements D. I. Mendeleev has the highest distribution. A short-range table is obtained from a long-period, if the 4th, 5th, 6th and 7th periods are cut to the elements of the IB group, push the resulting rows to be folded as before we have folded periods. The result is shown in Figure 6.24.

Lantanoids and actinoids here are also placed under the main table.

AT groupsthis table collected elements, at the atoms of which same number of valence electronsregardless of which orbitals these electrons are located. So, chlorine elements (typical element forming nonmetall; 3 s. 2 3p. 5) and manganese (element forming a metal; 4 s. 2 3d. 5), not possessing the selection of electronic shells, go here in the same seventh group. The need to distinguish such elements makes it allocate in groups subgroups: main- Analogs of A-groups of a long-range table and side - Analogues of B-groups. In Figure 34, the characters of the elements of the main subgroups are shifted to the left, and elements of side subgroups - to the right.
True, this arrangement of elements in the table has its advantages, because it is precisely the number of valence electrons first of all, the valence possibilities of the atom are determined.
The long-range table reflects the patterns of the electronic structure of atoms, the similarity and patterns of changing the properties of simple substances and compounds by groups of elements, a natural change in a number of physical quantities characterizing atoms, simple substances and connections throughout the element system and much more. A short-range table in this regard is less convenient.

Short-period table, main subgroups, side subgroups.
1. We convert a long-range table in the short-period element constructed from a natural range. Swipe the conversion.
2. Can it make a total valence electronic formula of atoms of elements of one group of short-period table? Why?

There are no clear boundaries at the atom. What is considered the size of an isolated atom? The atom core is surrounded by an electronic shell, and the shell consists of electronic clouds. EO size is characterized by a radius r. EO. All clouds of the outer layer have about the same radius. Consequently, the size of the atom can be characterized by this radius. It is called orbital radius atom(r. 0).

The values \u200b\u200bof the orbital radii atoms are shown in Appendix 5.
The radius of the EO depends on the charge of the core and on which orbital is an electron is located forming this cloud. Consequently, the orbital radius of the atom depends on the same characteristics.
Consider electronic shells of hydrogen and helium atoms. And in the hydrogen atom, and in the Helium atom electrons are located on 1 s.-Ao, and their clouds would have the same dimensions if the charges of the nuclei of these atoms were the same. But the charge of the kernel of the helium atom is twice as much as the charge of the nucleus of the hydrogen atom. According to the law of the Cool, the force of attraction, acting on each of the electrons of the helium atom, is twice as the force of attraction of an electron to the kernel of the hydrogen atom. Consequently, the radius of the helium atom should be much less than the radius of the hydrogen atom. And there is: r. 0 (He) / r. 0 (H) \u003d 0.291 E / 0.529 E 0.55.
At the lithium atom, the external electron is located on 2 s.-Ao, that is, forms a cloud of the second layer. Naturally, its radius should be more. Really: r. 0 (Li) \u003d 1,586 E.
At the atoms of the remaining elements of the second period external electrons (and 2 s., and 2 p.) Are accommodated in the same second electron layer, and the charge of the nucleus in these atoms with an increase in the sequence number increases. Electrons are stronger to the kernel, and naturally, the radii of atoms are reduced. We could repeat these reasoning and for atoms of the elements of the remaining periods, but with one refinement: the orbital radius is monotonously decreased when it is filling out each of the sublevels.
But if you distract from particularities, then the overall 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 orbital atoms decrease, and in the group - increase. The largest atom is the cesium atom, and the smallest atom of helium, but from the atoms of the elements forming chemical compounds (helium and neon they do not form), the smallest is the fluorine atom.
In the majority of atoms of elements in a natural row after lanthanides, orbital radii is somewhat less than expected, relying on common patterns. This is due to the fact that 14 lanthanides are located between lanthania and hafnium in the system of elements, and, therefore, the charge of the hafnium atom at 14 e. More than Lantana. Therefore, the external electrons of these atoms are attracted to the kernel more than they would be attracted in the absence of lanthanides (this effect is often called "lantanoid compression").
Note that when moving from atoms of elements VIIIA groups to atoms of elements IA groups, the orbital radius increases jumps like. Consequently, our selection of the first elements of each period (see § 7) turned out to be correct.

Orbital radius of an atom, its change in the system of elements.
1. From the data given in Appendix 5, build a graph of the orbital radius of an atom from the sequence number of the element for elements with Z. From 1 to 40. The length of the horizontal axis 200 mm, the length of the vertical axis is 100 mm.
2. How can I characterize the view of the resulting broken line?

If you inform the electron in an additional energy atom (as it can be done, you will learn from the course of physics), then the electron can go to another JSC, that is, the atom will be in excited state. This state is unstable, and the electron will almost immediately return to its initial state, and excessive energy is released. But if the energy reported by an electron is sufficiently large, the electron can completely break away from the atom, at the same time ionized, that is, turns into a positively charged ion ( cation). The energy required for this is called atom ionization energy(E. and).

Tear the electron from the only atom and measure the energy necessary for this is quite difficult, therefore practically determined and use molar energy ionization(E and M).

The molar energy of ionization shows what the smallest energy that is necessary for the separation of 1 mole of electrons from 1 praying atoms (one electron from each atom). This value is usually measured in kilodzhoules on mole. The molar energy of the ionization of the first electron for most elements is given in Appendix 6.
How does the atomic ionization energy 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 meaning, the ionization energy is equal to the work that needs to be spent on overcoming the force of attraction of an electron to the atom when the electron is moved from an atom to an infinite distance from it.

where q. - Electron charge, Q. - the charge of the cation remaining after removing the electron, and r. O is an orbital atom radius.

AND q., I. Q. - permanent values, and we can conclude that, work on the separation of an electron BUT, and with it and the energy of ionization E. And, inversely proportional to the orbital radius of the atom.
After analyzing the values \u200b\u200bof the orbital radii atoms of various elements and the corresponding ionization energy values \u200b\u200bgiven in Appendices 5 and 6, you can make sure that the relationship between these values \u200b\u200bis close to proportional, but is different from it. The reason that our conclusion is not very good with the experimental data is that we used a very rough model that does not take into account many essential factors. But even this rough model allowed us to make the right conclusion that with an increase in the orbital radius, the atomic ionization energy decreases and, on the contrary, with a decrease in the radius - increases.
Since in the period with an increase in the orders of the orbital radius of atoms decreases, then the ionization energy is increasing. In the group, with an increase in the ordinal number, the orbital atoms radius, as a rule, increases, and ionization energy decreases. The greatest molar energy of ionization is the smallest atoms, helium atoms (2372 kJ / mol), and from atoms capable of forming chemical bonds - at the fluorine atoms (1681 kJ / mol). The smallest - in the largest atoms, cesium atoms (376 kJ / mol). In the system of elements, the direction of increasing ionization energy can be schedically represented:

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

Excited state, ionization, cation, ionization energy, ionization energy, changing ionization energy in the system of elements.
1. Using the data given in Appendix 6, determine which energy should be expeked to tear one electron from all sodium atoms with a total weight of 1 g.
2. Using the data given in Appendix 6, determine how many times more energy should be spent for separation by one electron from all sodium atoms weighing 3 g than from all potassium atoms of the same mass. Why is this attitude different from the relationship of the molar energies of the ionization of these same atoms?
3. According to the data given in Appendix 6, build a graph of the dependence of the molar energy of ionization from the sequence number for items with Z. From 1 to 40. The dimensions of the graph are the same as in the task to the previous paragraph. Make sure this schedule complies with the choice of "periods" of the element system.

The second most important energy characteristics of the atom - electron affinity energy(E. with).

In practice, as in the case of ionization energy, a corresponding molar value is usually used - molar Energy Energy().

The molar energy of the gear affinity shows what the energy released when the electrons is connected to one pole of neutral atoms (one electron to each atom). Like the molar energy of ionization, this value is also measured in kilodzhoules on mole.
At first glance, it may seem that the energy should not be released, because the atom is a neutral particle, and there are no electrostatic forces between a neutral atom and a negatively charged electron. On the contrary, approaching the atom, the electron, it would seem, should be repelled from the same negatively charged electrons forming the electronic shell. In fact this is not true. Remember whether you ever deal with atomic chlorine. Of course not. After all, it exists only at very high temperatures. Almost does not occur in nature even more stable molecular chlorine - if necessary, it has to be obtained using chemical reactions. And with sodium chloride (cooking salt) you have to deal constantly. After all, the cook salt every day is consumed by a person with food. And in nature it is found quite often. But after all, the composition of the table salt includes chloride ions, that is, chlorine atoms, which attached one "excess" electron. One of the reasons for this such prevalence of chloride ions is that chlorine atoms have a tendency to connect electrons, that is, the formation of chloride ions from chlorine atoms and electrons is distinguished by energy.
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 electronic sheath of the chlorine atom when switching to one-contact anionu. At the same time, how do you remember, energy 3 p.-Production decreases. There are other more complex reasons.
Due to the fact that several factors affect the value of the energy of the germination of an electron, the nature of changes in this value in the system of elements is much more complex than the nature of the change in ionization energy. In this you can verify, after analyzing the table shown in Appendix 7. But since the value of this value is determined, first of all, by the same electrostatic interaction as the value of the ionization energy, the change in it in the system of elements (at least in Groups) in general terms with a change in the energy of ionization, that is, the energy of affinity to the electron in the group is reduced, and in the period - increases. It is maximum at the fluorine atoms (328 kJ / mol) and chlorine (349 kJ / mol). The nature of the change in the energy of an electron affinity in the system of elements resembles the nature of the change of ionization energy, that is, the direction of increasing energy of the electron affinity can be schematically replicated like this:

2. In the same scale on the horizontal axis as in previous tasks, build a graph of the dependence of the molar energy of the gear affinity from the sequence number for the atoms of elements with Z. From 1 to 40 using Appendix 7.
3. How does physical meaning have negative energy of an electron affinity?
4.What of all atoms of the elements of the 2nd period, the negative values \u200b\u200bof the molar energy of the gear affinity have only beryllium, nitrogen and neon?

You already know that the tendency of the atom to give his own and attach other electrons depends on its energy characteristics (the energy of the ionization and energy of the germination of the electron). What atoms are more inclined to give their electrons, and what - to accept other people?
To answer this question, we will bring in table 15 all that we know about changing these inconsistencies in the system of elements.

Table 15. Changes in the tendency of atoms to the return of their and attaching other people's electrons