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Home  / GENERAL CHEMISTRY Textbook / Chapter 6. Molecule structure / FIEs of element and bonding energy

FIEs of element and bonding energy

A comparison of the data in respect to the number of electrons in the outermost shell, with the number of chemical bonds the given atom can form, has shown that the main precepts for chemical bond formation set for hydrogen molecule bond formation, are valid for other atoms as well. That is because the bond is of electric nature and is formed at the expense of two electrons (one from each atom). Because of this precept we should expect a correlation between the FIEs of the atoms and the bonding energies of dual-atomic molecules.

At present the experimental bonding energies (those that are necessary for breaking a 1-mole bond of a dual-atomic molecule in the gas phase) have been identified.

Experimental data on bonding energy for a number of dual-atomic molecules (in the gas phase) formed of atoms of the 2nd and 3rd periods are given in Table 6.4-1 and Figure 6.4-1.  

TABLE  6.4-1

Molecule A2

FIE of A (eV)

Bonding Energy (kJ/mol)

Molecule

A 2

FIE of A       (eV)

Bonding Energy (kJ/mol)

Li2

5.4

110

Na2

5.1

72

Be2

9.3

30

Mg2

7.6

8.5

B2

8.3

274

Al2

6.0

168

C2

11.3

602

Si2

8.1

314

N2

14.5

941

P2

10.5

477

O2

13.6

493

S2

8.1

421

F2

17.4

140

Cl2

13.0

240


6_4_1

             Figure 6.4-1 

Experimental data given in Table 6.4-1 and Fig. 6.4-1, show that the bonding energy between the atoms practically does not depend on the FIEs of the atoms to be bonded.

Thus, for example, the bonding energy of a molecule composed of two nitrogen atoms (N2) is equal to 941 kJ/mol (the FIE of N is equal to 1,406 kJ/mol) while the bonding energy in a fluorine (F2) molecule is equal to 140 kJ/mol, and the FIE of the F atom (FIEF) is equal to 1,682 kJ/mol which insignificantly differs from the FIE of a nitrogen (N) atom.

Analogously, the FIE of beryllium (Be), equal to 900 kJ/mol, has a value close to that of boron (B). FIEB = 800 kJ/mol. The same is true of the FIEC of carbon, which is equal to 1,088 kJ/mol. While the bonding energies in molecules Be2, B2, and C2 comprise 30 kJ/mol, 225 kJ/mol, and 602 kJ/mol respectively. That is, a molecule, made up of 2 beryllium  atoms, does not exist in the gas phase, it is unstable, while molecule C2 does not break up into atoms even at a temperature over 5,000ºC.

On the basis of the analyses given in Table. 6.4-1 one might think that when switching from a hydrogen molecule bond to dual-atomic molecules with more than one electron, there is either a principle difference (additional non-electric forces) or action of some additional factors that explain the above-mentioned inconsistencies like the absence of correlations between the FIEs of the atoms being bonded and the bonding energy - in accordance with the electric nature of bonding.

Explanations based on suppositions of new (non-electric) forces are not logic explanations. After such explanations, there is always the problem of how to explain the physical essence of these forces.

This is why we will begin by looking for additional factors that explain the absence of the expected correlations and the independence of experimental data relative to the FIEs of bonding energy in dual-atom molecules in the atoms being bonded. 

      Table (6.4-1) can be conditionally divided into four groups: 

Group I includes molecules composed of identical atoms whose bonding energies are below 40 kJ/mol. These molecules fall apart into their constituent atoms when in a gas phase.  

Group II includes dual-atomic molecules composed of identical atoms whose bonding energies vary between 400 kJ/mol and 1,000 kJ/mol. Indeed, the bonding energy in these molecules greatly differs towards the plus side as compared to the bonding energy of a hydrogen molecule whose energy is equal to 429 kJ/mol. 

Group III includes dual-atomic molecules composed of different atoms whose bonding energies range from 340 kJ/mol to 550 kJ/mol. 

Group IV includes dual-atomic molecules built of identical atoms whose bonding energies vary from 50 kJ/mol to 350 kJ/mol. 

Before we begin the explanation, let's specify the questions that we are to cover. The gathered experimental data on bonding energies between atoms of the 2nd and 3rd periods of the table of elements allows us to formulate the first question as follows: 

Why is it that the chemical bonding energy between the multielectron atoms is much smaller or much greater (Table 6.4-1) than in a hydrogen molecule (H2) ?

      TABLE  6.4-2 

          Chemical Bonding Energy (kJ/mol) in a Row of Dual-Atomic Molecules   

MOLECULES  AND  THEIR  BONDING  ENERGIES 
Molecule  Group I  Bonding Energy Molecule  Group II  Bonding Energy 
Be2 - 30 C2 - 602
Ne2 - 4 N2 - 941
Mg2 - 7.6 O2 - 493
Ar2 - 7 P2 - 477
S2 - 421
                                        

Molecule

Group III 

Bonding Energy Molecule

Group IV

Bonding Energy 
LiF - 572 B2 - 274
NaF - 447 Br2 - 190
LiCl - 480 Cl2 - 239
NaCl - 439 F2 - 139
Li2 - 110
Na2 - 72

 

The data on the number of electrons in the outermost shell allows us to ifentify the maximal number of electrons that can be situated in the outermost shells of the atoms of the 2nd and 3rd periods. The data on the FIEs of the atoms of these periods are given in Table 6.4-2.   

The following table will help us answer the question:

Why does the bonding energy in dual-atomic molecules differ so much in reference to the bonding energy in a hydrogen molecule? 

This table offers data on the atoms' affinities to the electron, which is also of great use for additional confirmation of the correctness of the conclusions made on the basis of the comparison of the FIE data. The data on the affinity of atoms to the electron indicates the energy that is discharged when the electron is bonded to the atom.

      TABLE   6.4-3 

      First Ionization Energies (FIEs) and Affinity to the Electrons in Elements  of  Periods 1, 2, and 3  in  the Table  of  Elements  in  kJ/mol 

            1        2       3       4       5        6       7        8        9       10

           H       He     Li     Be     B       C       N       O       F       Ne  

    FIE    1310  2372   519  900   799   1086  1406  1314  1682  208

    Aff.     67,4     <0      77    <0    31.8  119.7   4.5   141.8   349    <0 

 


 

           11      12      13     14      15      16      17      18

            Na     Mg     Al     Si        P        S       Cl      Ar

     FIE     498    736    577   787   1063  1000   1255  1519

     Aff.     117.2   <0      50    138     75    199.6   356     <0

 The affinity data given in the table is actually the result of studies of FIE values of atoms to which electrons were bonded. That is, for example,* the energy discharged when connecting the electron to the Na atom [(the affinity value for sodium (Na)] via the reaction      

 

      Na + e- → Na-1 

is measure by defining the FIE in reaction:  

      Na -1 → Na + e- 

The affinity values smaller than zero, indicated in the table, show that the electron does not bond to atoms with an affinity equal to <0.

Another question might arise:

What is common between the FIE value, the affinity of the atom to the electron, and the bonding energy? Or: Why do we think that the FIE values and those of the affinity can help us answer the question about the cause of the great difference in bonding energy in a number of dual-atomic molecules.

As we indicated in section 4, in the case of multi-electronic atoms during bond formation, just as in the case of a hydrogen molecule, chemical bonding takes place via the electronic pair (one from each of the atoms being bonded) while both bonding electrons enter the outermost shell of both atoms to be bonded.

The observance of this rule for multi-electronic atoms was proven on the basis of chemical experimental data relative to the existence of stable compounds (for example, compounds of atoms of the 2nd and 3rd periods were taken) where the number of electrons surrounding each atom and entering the molecule, does not exceed eight. This value coincides with the maximal number of electrons that can exist in the outermost shells of the atoms of the 2nd and 3rd periods.

To explain the significant deviation of the bonding energies in multi-electronic atoms, as compared to that in a hydrogen molecule, it is necessary to deepen our understanding of the reason why the number of electrons in the outermost shell is limited.

It was already shown that the main forces in the atomic and molecular systems are the electric forces that attract the electrons to the nucleus and the interelectronic repulsion forces.

The increase of the attraction forces increases the absolute value of the potential energy of the system: neutral atom + free electron. On the other hand, the increase of the interelectronic repulsion forces decreases the absolute value of the system's potential energy: neutral atom + free electron.

The bonding of the electron to the atom occurs when there is energy gain, or, in other words, if the absolute value of the potential energy of the system atom + electron increases as a result of the bonding of the electron to the atom. The data on the affinity of the atom to the electron shown in table 6.4-3 gives us a digital value of the energy gain during the bonding of the electron to the atom.

During the bonding of the electron to the atom, the total attraction energy of the electrons to the nucleus increases at the expense of the increase of the number of electrons attracted to the nucleus. On the other hand, the inter-electronic repulsion energy increases because of the increase of the amount of electrons. That is, the bonding of the electron to the atom takes place if, as a result of this bonding, the attraction energy gain is greater than the energy loss received at the expense of the increase of the repulsion energy.

To illustrate this, let's have a look at the inter-electronic repulsion forces, the electron-nuclear attraction, and the effective nuclear charge when the electron is attached to a hydrogen-like atom (composed of one nucleus and one electron) and a helium-like atom (composed of one nucleus and two electrons).

When an electron is bonded to a hydrogen atom, there appears an inter-electronic repulsion force, and the electrons' attraction energy to the nucleus undergoes a change.

Let's calculate the potential energy change that occurs during the bonding of an electron to a hydrogen atom. The hydrogen atom turns into a helium-like atom with a nuclear charge of 1 proton unit and with 2 electrons at equal distances on either side of the nucleus. The potential energy of this system is equal to the difference between the electrons' attraction energy to the nuclei and the repulsion energy.

In the cited case (when the nuclear charge and the electron's charge have identical absolute values), the electrons' attraction energy to the nucleus is equal to: 2e2/R, where R is the radius of the orbit along which the electrons rotate. The inter-electronic repulsion energy is equal to: e2/2R. The total potential energy is equal to the difference between the attraction energy and that of inter-electronic repulsion. That is:

  2e2 / R - e2 / 2R = 3e2/2R                   (6.4-1) 

In accordance with equation (6.1-1) the energy of a helium-like atom is equal to: E = [13.6 (z - 0.25)2] · 2. At z = 1 (cited example):

 E = [13.6 (1 - 0.25)2] · 2 = 15.3 eV.  

Here 13.6 eV is the hydrogen atoms' energy (EH).

According to equation (5.6):  

EH = e2 / 0.529 

where 0.529 is the radius of a hydrogen atom. The potential energy is equal to 2EH. Thus: 2e2 /0.529 = 27.2 eV and e2 = 14.38. By substituting e2 in equation 6.4-1, we get: 3 · 14.38/2R = 30.6 eV, — where 30.6 eV is the double energy (15.3 · 2) of the helium-like atom's charge 1.

 Thus: R = 0.705 Å 

Knowing the radius of the helium-like atom, we can calculate the change of both the potential attraction energy and the potential repulsion energy, since we know that (e2 / 0.529) · 2 = 13.6 eV (energy of 1 gram of hydrogen atoms): e2 = 14.4 eV.  From this we get:

2e2 / R = 2 · 14.4 / 0.705 = 40.8 eV; e2 / 2R = 10.2 eV.     

The potential energy of the system, where the electron is separated from the hydrogen atom at an infinite distance, is equal to the potential energy of a hydrogen atom:  2 · 13.6 = 27.2 eV.  

That is, when the electron is bonded to a hydrogen atom, the potential energy of the electrons' attractions to the nucleus increases by 13.6 eV. On the other hand, when bonding an additional electron to a hydrogen atom, the interelectronic repulsion energy increases by 10.2 eV. There is no inter-electronic repulsion in the initial hydrogen atom.

Thus, when adding an electron to a hydrogen atom, the absolute potential energy value increases by 3.4 eV (13.6 - 10.2 = 3.4). The potential energy of 1 gram of hydrogen atoms is equal to 27.2 eV (double value of the hydrogen atoms' energies).

Respectively, the energy of 1 gram of helium-like atoms, with a charge equal to that of a hydrogen atom, constitutes, according to the calculation: 30.6 eV (27.2 + 3.4 = 30.6). According to experimental data, the potential energy of such atoms comprises

      28.7 eV [(13.6 + 0.74) · 2 = 28.7, 

where 13.6 is the energy of a hydrogen atom; 0.74 eV is the affinity of the hydrogen atom to the electron. 

Processes proceeding via attraction forces (a stone falling onto the ground, an electron striving towards a nucleus, etc.) occur spontaneously with a discharge of energy. That is, in accordance with the calculation, the bonding of the electron to a hydrogen atom occurs spontaneously with a discharge of energy. In other words, the hydrogen atom should have a positive affinity to the electron. This has been confirmed via experiments.

Now let's cite the change of the potential energy when bonding an electron to a helium atom. The above-described calculations have shown that the change of the system's energy, during the bonding of the electron, is equal to the value of the potential energy change divided by 2. According to the virial theorem, half of the system's potential energy is equal to the system's energy. That is, when calculating the energy of a system consisting of electrons and nuclei, we need not calculate the kinetic energy of the electrons.

It is enough to define the potential energy of the system (i.e., the distance between the electrons and the nuclei) and divide the result by two. To compare the calculated and experimental results, we define the ionization energies of all the electrons. The sum of all these energies, according to the above mentioned, is equal to the energy of the system.

According to equation (6.1-1) the energy of 1 gr. mol of atoms with a nuclear charge of 2 proton units and with 2 electrons orbiting in a circle around the nucleus, on one plane is calculated by the equation:

EHe = 13.6 (2 - 0.25)2 · 2, 

where 13.6 is the energy of the hydrogen atom; 2 is the nuclear charge of the helium atom; 0.25 is the allowance for inter-electronic repulsion.

Analogously, the energy of 1 mol of atoms with a nuclear charge of 2 proton units and with 3 electrons rotating around the nucleus in one circle can be calculated by the equation: 

      E3 = 13.6 (2 - 0.577)2 · 3, 

where 0.577 is the inter-electronic repulsion allowance (relative to three electrons). See the book: "How Chemical Bonds Form and Chemical Reactions Proceed" (p.42). The calculation concludes as follows: 

 E2 = 83.3 eV ;           E3 = 82.6 eV  

That is, the bonding of the electron leads not to the increase of the potential energy, but to its decrease. Recall the fact that the absolute value of the system's energy is equal to half of its potential energy. The bonding of one electron to a helium atom is accompanied by a decrease of the absolute value of the potential energy, and therefore, cannot take place spontaneously, i.e., without the use of energy. Indeed, stones do not reveal themselves from under the ground spontaneously.

In other words, the calculation shows that an electron cannot bond to a helium atom, which has been proven experimentally. The affinity of a helium atom is less than zero. 

Therefore the bonding or nonbonding capacity of the electron to the atom is defined by the difference in the change of the absolute values of the potential attraction energies of all the electrons to the nucleus and the mutual interelectronic repulsion. If this difference is greater than zero— the electron will bond; if it is smaller than zero — it won't.

The data on atoms' affinities to the electron given in Table 6.4-3 show that for the atoms of periods 1, 2, and 3 in the table of elements, besides the atoms of helium (He), beryllium (Be), magnesium (Mg), neon (Ne), and argon (Ar), the atoms' affinities to the electron are greater than zero. During the bonding of electrons to atoms of periods 2 and 3 (besides Be, Mg, Ne, Ar) there is an energy gain. Thus, the increase of the attraction energy during the bonding of electrons to the nucleus is greater than the increase of the repulsion energy.  

In the case of He, Be, Mg, Ne, and Ar atoms, the increase of the attraction energy during the bonding of the electrons to the nucleus, is smaller than the energy increase of the inter-electronic repulsion. An independent confirmation of this conclusion is the data on the FIEs for atoms of the 2nd and 3rd periods given in Table 6.4-2. Each of the numbered elements differs from the next by nuclear charge and by the amount of electrons surrounding the nucleus. 

The atom of each of the following elements has a positively charged nucleus by one proton more. The number of electrons in the electronic shell of each of the following atoms is by one electron more than in the previous shell.

The transition from one element to the next, for example, from sodium (Na) to magnesium (Mg) can be represented schematically as two consecutive processes: first the nuclear charge of the Na atom in-creases by one proton unit and turns into a Mg nucleus, then one electron is bonded to the atom that has a nuclear charge of 12 proton units and an electronic shell that has 11 electrons [shell of a sodium (Na) atom].

With such consideration, the atoms' FIE values correspond to the energy gain during the bonding of the electron to an atom whose nuclear charge had been increased by one proton unit.

The FIE values of Table 6.4-3 increase the values of the atoms' affinities to the electrons. For example, the FIE for Na is equal to 498 kJ/mol, while the affinity is equal to 117.2 kJ/mol. Thus, the bonding of an electron to an atom with a nuclear charge of 11 proton units and surrounded by 10 electrons offers energy gain of 498 kJ/mol. The bonding of an electron to an atom with a nuclear charge of 11 proton units and surrounded by 11 electrons offers energy gain of about 4 times smaller (117.2 kJ/mol). That is, during the bonding of an electron to an atom, the increase of the nuclear charge abruptly increases the energy gain.

During chemical bond formation, the number of electrons in the atoms' outermost electronic shells increases by one electron, and the effective charges of the atoms being bonded are changed. The effective charges of the nuclei to be bonded are changed because of the attraction of the charged nuclei and because of the increase of the number of electrons in the outermost shells of the atoms being bonded.

In order to compare the bond formation process with the processes of bonding an electron to an atom without changing the nuclear charge of the atom, and the process of bonding the electron to the atom with a simultaneous increase of its nuclear charge by one proton unit, we should evaluate the influence of the atoms being bonded on the effective nuclear charge, and the attraction of the nuclei of these atoms that occurs during bond formation.

An additional attraction force of the electron to the atom's nucleus in a hydrogen molecule (force occurring as a result of the attraction of the atom's nucleus) is equal to the projection of force F1 that attracts the electron to a nucleus that bonds the same electron to another nucleus. (See Figure 6.1.1) The value of this projection force is equal to:

  F1 · cos 60º =  0.5 F1.    

That is, the mutual approach of the nuclei leads to the increase of the attraction force of the bonding electrons to the nuclei by 50%, which is equal to the increase of the effective charge of the nuclei to be bonded by 0.5 proton units.

Bond formation is, from the viewpoint of the energy gain, a sort of middling process between the bonding of the electron to a neutral atom (measured affinity to the electron) and the bonding of the electron to the atom whose nuclear charge is increased by one unit.

      The data on the affinity of atoms to the electron and the data on the FIEs allow us to elucidate exactly why the bonding energy in molecules given in table 6.4- 2 is much smaller than the bonding energy in a hydrogen molecule.

We can say that when an additional electron is introduced to the outermost shells of Be and Mg, the inter-electronic repulsion energy is increased to a greater extent than is the attraction energy of the electrons to the nucleus

According to the data in Table 6.4-3, when going from lithium (FIE - 519 kJ/mol) to beryllium (FIE - 900 kJ/mol), the FIE increases by 400 kJ/mol; but when going from beryllium to boron (FIE - 799 kJ/mol) the energy gain decreases to 100 kJ/mol.

According to Table 6.4-3, there are 3 electrons in the outermost electronic shell of boron, while there are 2 electrons in the outermost shell of beryllium. That is, when the electron is bonded to beryllium with a simultaneous increase of the nuclear charge by one proton unit, the electron being bonded enters the existing outermost shell of beryllium, and the energy gain will thus be by 100 kJ/mol smaller than during the entrance of the electron to the outermost shell of lithium (when going from lithium to beryllium).

Thus the decrease in the energy gain, when the electron enters the outermost shell of a beryllium atom, can be evaluated as greater than 100 kJ/mol and smaller than 400 kJ/mol.    

During the formation of molecules Li2, Be2, B2, two electrons enter the outermost shells of the atoms being bonded (one electron into each atom). In the case of lithium and boron, according to the data on the affinity of atoms to the electron, the energy gain amounts to 77 kJ/mol and 32 kJ/mol respectively for each of the two atoms being bonded, and therefore these atoms have 154 kJ/mol and 64 kJ/mol respectively.

In the case of beryllium, when the electron enters the outermost shell of this element, according to the atoms' affinity to the electron, energy is not gained, but even lost. According to the above - the energy loss is generally equal to 100 kJ/mol or 200 kJ/mol.

Since the energy gain, say, during hydrogen molecule formation out of hydrogen atoms with a positive (>0) affinity value to the electron, comprises 250 kJ/mol, the abrupt decrease of bonding energy for atoms with a negative atom affinity to the electron, mentioned in Table 6.4-3, is quite comprehensible.

The explanation concerning the anomalously small bonding energies in elements of  GroupsII and VIII of the table of elements is an independent semi-quantitative proof of the fact that during covalent bond formation, both bonding electrons enter the outermost shells of the atoms to be bonded. This conclusion was made only on the basis of the comparison of the data of the number of electrons in the outermost shell with the elements' valences.

The fact that atoms of Group II form stronger bonds, as compared to the inert gases, is also proof that in the course of covalent bond formation, the effective charge of the atoms to be bonded increases. 

Now let's answer the question: Why is the bonding energy in dual-atomic molecules shown in Table 6.4-1, namely, in molecules of carbon (C2), nitrogen (N2), oxygen (O2), etc. 1.5 to 2 times greater than the bonding energy in a hydrogen molecule?

The outermost shells of carbon (C), nitrogen (N), and oxygen (O) atoms contain 4, 5, and 6 electrons respectively. The number of bonds these atoms form is limited by the number of additional electrons that enter their outermost shells. Thus, atoms of C, N, and O can form 4, 3, and 2 chemical bonds respectively. And not one but several chemical bonds can be formed between any two atoms listed in Table 6.4-2, which presupposes a much greater energy gain, as compared with the formation of 1 bond in the case of a dual-atomic molecule where the atoms being bonded each have 1 electron in the outermost shell.

Additional confirmation of the correctness of this explanation are the data on the bonding energy in multi-atomic molecules, where the possibilities for the formation of additional bonds between two atoms of C, N, and O come to naught at the expense of bond formation, say, with hydrogen atoms.

It was found experimentally that the bonding energy between carbon atoms in molecule H3C - CH3, nitrogen atoms in molecule H2N - NH2, and oxygen atoms in molecule HO - OH is by 1.5 to 2 times smaller and close in value to the bonding energy of a hydrogen molecule.  

When atoms are bonded with one chemical bond, such a bond is known as a common chemical bond or a single chemical bond. When atoms are bonded with several common or single bonds, thereby forming double or triple ones, such bonds are known as multiple bonds.

Single bonds, when describing chemical structures, are illustrated with a dash (—). For example, the structure of methane (CH4) is described by the formula: 

H
Ι
H C H
Ι 
H

Each dash represents a pair of bonding electrons. The electronic structure of methane is described by the formula:  

      H

      H  : C :  H

      H 

Here the dots represent bonding electrons. Multiple bonds, for example, in nitrogen (N2) and oxygen (O2) molecules are described by structural formulas:

      N  Ξ  N    and   O   =   O

and   by electronic ones:       

. . . .
: N : : : N : and : О : : О :
. . . .

Here the electrons between the atoms are bonding ones, while those, sort of on the outside, are nonbonding ones (which do not take part in bond formation).  

According to electronic formulas N2 and O2 in nitrogen (N), after the formation of a triple chemical bond, there remain 2 nonbonding electrons or 1 free electronic pair with each of the atoms. After the formation of a dual chemical bond in oxygen, 2 free electronic pairs remain in each of the 2 atoms bonded via a dual bond in molecule O2.

Now let's take the question:  Why is the bonding energy in some of the dual-atomic molecules, which are made up of different atoms, much greater than that of other molecules that are made up of identical atoms?

According to experimental data (Table 6.4-2.) the bonding energies in molecules Na2 and Cl2 comprise 74 kJ/mol and 242kJ/mol respectively, while the bonding energy in molecule NaCl is equal to 439 kJ/mol. In all the three molecules (Na2, Cl2, and NaCl) the bonding takes place at the expense of the attraction of the atoms' nuclei in sodium and chlorine to the electronic pair rotating in a plane perpendicular to the axis connecting the atoms' nuclei.

From this point of view, it seems logic that the total attraction energy of the atoms' nuclei in the Na and Cl to the bonding electronic pair should have an average value between 74 kJ/mol and 242 kJ/mol (the attraction energies of two Na atoms' nuclei and two Cl atoms' nuclei, respectively, to the bonding pair of electrons).

Sodium and chlorine atoms differ greatly relative to their affinity to the electron. As previously indicated, bond formation can be represented as a dual process on the first stage of which the energy gain is received at the expense of the atoms' affinities to the electrons. That is, from this point of view, the energy gain, when forming a Cl2 molecule, should be greater than when forming a NaCl molecule.    

Chapter 6. Molecule structure >> 
Conclusions >>   
**Molecules formed of multi-electron atoms >>  
**Ionization energy of multi-electron atoms >>       
**Chemical Energy. FIEs of element and bonding energy
**Chemical Bonding Energy >>
***Bond Lengths >>
Valence >>
Conclusions >>     
**Donor - acceptor Bond (DAB) >>   
Van der Waals Bond (VWB) >>    
Dynamic Bonds >>
Conclusions >> 
Chapter 6 Textbook Questions