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Home  / GENERAL CHEMISTRY Textbook / Chapter 6. Molecule structure / Dynamic Bonds

Dynamic Bonds

Studies conducted on energies and bond lengths in chemical compounds, where one atom is bonded with various types of bonds to other atoms, have shown that, as a rule, the energy of weak bonds (DABs and VWBs) increases while that of the strong bonds (homopolar and heteropolar covalent bonds) decreases.

Simultaneously the length of the weak bonds decreases, while that of the strong ones increases. This fact was realized only after the studies of the phenomenon called electron-nuclear isomerization, which will further be referred to as electronic isomerization or electron-nuclear isomerization.

It has been found experimentally (see details Appendix A 16.), that the electrons in molecules transit so rapidly, that they cause isomerization of the chemical compound. Thus, for example, if we take a cation like ammonia (NH4)+ the process of electronic isomerization can be demonstrated thus: 

H N : H ⇆  H : N H

Here the two dots (:) indicate covalent bonds; the single arrows (← →) indicate DABs. Since the speed of the electronic transition by far exceeds that of the nuclei, the hydrogen atoms' nuclei occupy an intermediate position. That is, the bond lengths between all the hydrogen atoms and the nitrogen atom are the same, but greater than the covalent bond and smaller than the DAB.

The thermal stability of the compound (at high temperature) is defined by the energy of the weakest bond in the compound. That is why electronic isomerization leads to the increase of the thermal stability in the compound.

The greatest effects of isomerization (including the compounds' stability increase) are evident in cases when, as a result of this isomerization, identical electronic isomers are formed. In the above cited example, the electronic isomerization of an ammonia ion (NH4)+ has an identical initial and final molecule and has 3 covalent bonds of N and H and one DAB of nitrogen-hydrogen.

That is, the electronic energies of both isomers are the same, and the electronic isomerization leads to a complete equalization of both the lengths and energies of all the bonds in ionic ammonia (NH4).

Cases when different bonds bond one atom to other atoms are a common phenomenon. Thus, for example,

in molecule   

N = = О

the nitrogen atom is bonded with two oxygen atoms via double and single bonds.

Analogously, in a benzene molecule:                          

The carbon atoms are bonded between themselves via single and double covalent bonds. In the anion NO3 the oxygen and nitrogen atoms are bonded via single and double covalent bonds and DABs: 



O = = N O

Likewise are bonded the oxygen anion (O) together with the chlorine in the anions ClO2- (chlorite ion), ClO3- (chlorate ion), ClO4- (perchlorate ion), covalent bonds, and DABs.  In all the above-mentioned compounds, as a result of electronic isomerization, we see the equalizing of various bonds in respect to their lengths and energies.

Electronic isomerization allows us to realize why substances, composed of molecules, and having, according to the theory, a DAB or a VWB, are stable at temperatures much higher than those necessary for breaking these bonds. For example, isomerization explains the existence of such compounds as (I-I-I)-, (Br-Br-Br)(Cl-I-Br)-.

In such molecules anions of halogens (I-, Br-, Cl-) are bonded strongly with molecules of I2, Br2, IBr. These have 8 electrons in the outermost shell; that is, this shell is filled. Halogen atoms in halogen molecules just like I2, Br2, IBr, also have filled electronic shells. That is why there can be only a VWB bond between a halogen anion and a halogen molecule, which is a very weak bond that excludes the synthesis of a substance like I3- , Br3-, Cl3-.

The bonding energy value between a halogen anion and a halogen molecule is close to that of a VWB formed between the atoms of inert gases; that is, about 10 kJ/mol. The bond's length is just over 4Å.

It has been found experimentally that compounds like  I3-, Br3-, Cl3-are stable. In these compounds the length of the bonds of the topmost atoms with a central I atom comprise 2.9Å for I - I - I, and 2.53Å for Br - Br - Br. These values are close to those of a common covalent bond in molecules I2 and Br2, which are equal to 2.66Å and 2.28Å respectively.

According to the structure of these compounds, there should proceed electronic isomerization along the scheme:

  I- ... I - I      ⇆     I - I ... I-

Dots (...) indicate VWBs; the dash (-) indicates covalent bonds.

As a result of the electronic isomerization, we get a compound where both I - I bonds become equal in energy and length.

The structure of stable compounds of inert gases with halogens such as XeF2 is explained identically as that of I3. The structural formula of such a compound can be demonstrated analogously to that of I3- thus:                              (F : Xe)+ ... :F-

Here (:) indicates a covalent bond formed at the expense of one electron of a fluorine (F) atom and one electron of a xenon (Xe) atom. The three dots (...) indicate a VWB which is formed between the fluorine (F) anion and the cation (F : Xe)+. Therefore, this VWB should be much stronger than a common VWB, at the expense of the electrostatic attraction.

There is an additional contribution to the strengthening of this bond offered by the process of electronic isomerization, which proceeds via the following scheme:

(F : Xe)+... F- ⇆  F- ... (Xe : F)+

The outermost shells of all the atoms in molecule (F : Xe)+...F- are completely filled (there are 8 electrons). In order to have such a bond formed, we should compensate the energetic expenditures connected with the transition of the electrons from the xenon (Xe) shell to the fluorine (F) shell with the formation of anion-fluorine (F-).

The FIE of xenon is equal to 12.13 eV; the affinity of fluorine to the electron is equal to 3.34 eV. That is, the transition of the electron from the xenon to the fluorine requires the expenditure of a large amount of additional energy, i.e., 8.79 eV (12.13 - 3.34 = 8.79 eV). The compensation of this energy is possible at the expense of the addition of a second fluorine atom to the xenon with the formation of a covalent bond: xenon (Xe) - fluorine (F). Here the bonding pair of electrons, formed of single-fluorine and single-xenon electrons, enter the outermost shells of both fluorine and xenon.

The formation of the covalent bond xenon-fluorine in the course of the formation where the number of electrons in the outermost shell of xenon increases by 1, compensates the energetic expenditures needed for the transition of the electron from the xenon to the fluorine. In the course of the formation of this covalent bond, the electron transits from the fluorine atom to the cation-xenon (Xe+) shell, which presupposes the discharge of energy equal to that of the xenon's FIE (12,13 eV).

Thus, according to the theory, we should get a synthesis of stable compounds of xenon and fluorine only of types XeF2, XeF4, and XeF6. The synthesis of these compounds and the lack of literature describing such compounds as XeF, XeF3, and XeF5 serve as experimental confirmation of the correctness of our theoretical reasoning.

Now let's take some inert gas compounds and study possibility their synthesis.

What stable compounds can we expect on the basis of the theory of chemical bonding, and what stable compounds, according to the theory, are hardly possible?

The union of inert gases and halogens, generally, can be seen in the formula: In Haln (where In is the inert gas, Hal is the halogen, and n is the even number 2, 4, or 6).

In the general electronic formula, when n = 2 these compounds can be represented as Hal : In ... Hal (where (:) corresponds to the covalent bond and (...) to the VWB.

As a result of electronic isomerization, the energies and lengths of the bonds equalize. As the energy of the VWB is about ten times smaller than that of the covalent bond, we can suppose that the equalized energy of both bonds in the molecule, as a result of the isomerization, are defined by the energy of the covalent bond Hal : In.

Now let's see how the inert gas halogen depends on the type of halogen and type of inert gas according to the bonding energy theory.

Each group of halogens and inert gases has the same number of electrons in the outermost shell. The halogens have 7 electrons while the inert gases have 8. These atoms differ mainly by their FIEs.

The FIEs of the inert gases in the following row are:

He (24.58 eV), Ne (21.55 eV), Ar (15.76 eV),
Kr (14.00 eV),
Xe (12.13 eV), Rn (10.75 eV).

The FIEs of the halogens in the following row are:

F (17.42 eV), Cl (13.01 eV), Br (11.84 eV),
I (10.42 eV).

According to the general bonding theory, the energy of the homoatomic covalent bond, when the FIE is greater than 11 eV, will decrease when the FIE increases. According to experimental data, the bond in the fluorine molecule (F2) comprises 159 kJ/mol, while that of chlorine (Cl2) comprises 242 kJ/mol. The FIEs for F and Cl comprise 17.4 eV and 13.01 eV respectively. That is, the compound of the ArF2 type is less stable than compound KrCl2.   

According to our theory of chemical bonding, the energy of heteroatomic covalent chemical bonds is higher than that of homoatomic chemical bonds; and the energy of the bonds increases proportionally to the FIE differences of the atoms being bonded.

Therefore, the strongest heteroatomic covalent bonds can be expected thus: F : Rn and F : Xe since the FIE of fluoride (F) is equal to 17.4 eV, that of Xe - 12.13 eV, Rn - (10.74 eV). The energy of this bond is » 400 kJ/mol according to experimental data. The energies of the rest of the heteroatomic covalent bonds are much smaller.

Besides covalent bonding energy, another main condition for stable compounds of the halogen + inert gas type: Hal : In ... Hal, is that the electronic energy in these compounds should be smaller than the total of the electronic energies of the halogen molecules (Hal2) and the inert gas atoms (In), which can form out of them as a result of electronic isomerization via reaction:

Hal : In ... Hal → Hal2 + In

or:  F : Rn ... F→ F2 + Rn

Thus, the stability of the compounds of inert gases and halogens is defined as both covalent bond energy Hal : In and bonding energy Hal : Hal, formed via the breaking of this compound into a halogen molecule (Hal2) and an inert gas atom. The smaller is the bonding energy in the halogen molecule (Hal2), the greater is the stability of the compound halogen + inert gas which is described as Hal : In ... Hal.

Fluorine (F) has the smallest bonding energy (» 159 kJ/mol) in the halogen row. This element, as indicated above, offers a very strong bond with xenon (Xe) and radon (Rn) atoms. According to the above mentioned, the bonding energy between fluorine and xenon or radon is about 400 kJ/mol.

That is, molecule F : Xe ... F is stable. It will not change into a fluorine molecule and a xenon atom all by itself.

Likewise, in accordance with the bonding theory, the bonding energy of fluorine and xenon in molecule F Xe F can be evaluated at » 200 kJ/mol. According to experimental data, the energy of this bond is equal to 130 kJ/mol. Analogous studies have shown that in the cases of Cl : Kr ... Cl Cl: Xe ...Cl and  Cl : Rn ... Cl, the covalent bonding energy Cl : Xe amounts to » 240 kJ/mol while the bonding energy in a chlorine molecule comprises 240 kJ/mol. This presupposes both a significantly lower thermal stability of such compounds as Cl Xe Cl and a lower chemical stability. That is, it presupposes their transformation into a chlorine molecule and an atom of inert gas.

Thus, according to the above, we can expect only a synthesis of stable compounds of halogens with inert gases of type Fn Xe and Fn Rn, where n = 2, 4, 6. The synthesis of compound xenon-fluorine (XeF2 , XeF4 , XeF6) and the unsuccessful attempts at synthesizing ArF2 , KrCl2 , and XeBr2 are experimental  proofs  of  this  conclusion.

The elucidation of the comparatively exotic compound structure of inert gases with oxygen and halogens gives us to understand the structure of the widely used compounds in chemistry such as: sulfuric acid (H2SO4), phosphoric acid (H3PO4), phosphorus pentachloride (PCl5), and sulfurous hexafluoride (SF6).

In the framework of the explanations given in textbooks, it was supposed that the central atom in these compounds [sulfur (S) and phosphor (P)] contained more than 8 electrons in the outermost shell. It was also supposed that the extra electrons were located on the 3d orbitals.

This explained the cause of finding more than 8 electrons in the outermost shell, and the difference between phosphorus (P) and sulfur (S) on the one hand, and nitrogen (N) and oxygen (O) on the other, which do not form analogous compounds. Nitrogen (N) and oxygen (O), according to quantum chemistry, did not have any closely situated 3d orbitals, and correspondingly, did not have any free space for extra electrons.

This explanation was seriously contradicted both experimentally and by the quantum-chemical theory itself. According to the theory and experiment, the filling of the 3d orbitals by electrons began with scandium (Sc). Between Sc and S in the table of elements, there are four elements (Cl, Ar, K, Ca), which have no electrons on the 3d orbitals.

That is, if the 3d orbitals of the elements of the 3rd period were energetically available, then the chlorine (Cl) and argon (Ar), whose nuclear charges are greater than those of phosphorus and sulfur, would bond the extra electrons to the 3d orbitals. Relatively, these elements (Cl, Ar), seemingly, should have a valence greater than 1 (chlorine) and greater than zero (argon). This has never been proven experimentally.

How is the structure of these compounds explained in the framework of our new approach?

In complete analogy with compounds of inert gases with oxygen and halogens (compounds, phosphor and sulfur are also connected with either halogens or oxygen) the structure of sulfuric and phosphoric acids can be shown thus: 

HO S OH           HO P OH

Here there is a covalent bond between the hydroxyl oxygen and the central atom (sulfur and phosphor), and there is a DAB between the sulfur and phosphor with an atom of oxygen. This is a complete analogy of the bonding in molecules of the XeO3 type.

Compounds of phosphor (P) and sulfur (S) with halogens are structured analogously to XeF2. That is, phosphorous pentachloride is described by the following structural formula:

                                     has  four covalent bonds and two VWBs.                             

In all the described compounds the energies and the bond lengths between the central atom and the oxygen or halogen are equalized at the expense of electronic isomerization.

The similarity of bonds XeF4 and PCl5 is confirmed by the similarity of the types of formation and break-up reactions of these compounds. It is known that when heating five-chlorine phosphor, a reverse reaction takes place:

PCl5 ⇆ PCl3 + Cl2

Likewise, just as in xenon bonds with fluorine, in the case of sulfur and phosphorous well known stable compounds PCl3, PCl5, SF4, SF6 and the unknown PCl4, PCl6, SF3. SF5.

Analogous to the compounds of inert gases with oxygen and halogens, we can explain why nitrogen and oxygen do not form the same kind of compounds.

Stable compounds form only when the FIEs of the halogens and oxygen are greater than that of the inert gas, which is the central atom. In our case, the FIEs for sulfur (S) and phosphorus (P) are 10.36 eV and 10.48 eV respectively. The FIEs for oxygen (O), chlorine (Cl), and fluorine (F) are 13.61eV, 13.01eV, and 17.40 eV respectively. On the other hand, the FIE of nitrogen (N) is equal to 14.53 eV.

Remember that compounds of xenon (Xe) with a FIE of 12 eV, are the only stable compounds of inert gases with oxygen and halogens. Krypton (Kr) with a FIE of 14 eV, does not form such compounds.

Analogously to the above-cited compounds, we can explain the structure of compounds where the central atom is bonded to identical atoms via double and single bonds like SO2, SO3, NO2, C6H6, etc.

Experiments have proven that these compounds also reveal the tendency to equalize the lengths and energies of single and double bonds.

Up till now, school texts and scientific literature explain the equalizing of bonds, relative to their energy and length, in the framework of quantum chemistry of the resonance theory.

The resonance theory presupposes that such a structure, as SO2 is a super-position of two structures, which is illustrated thus: 


The possibility of electronic isomerization in the framework of the resonance theory, is not only rejected, but even ridiculed: "This is the kind of foolish propositions the chemists had before quantum chemistry was introduced."

In the framework of the above said, between the structures SO2 there is dynamic equilibrium:


which, in this case, as in all other cases, leads to the equalizing of chemical bonds. 

One very interesting example involves a combination of the VWB and polar bond for a hydrogen atom. In chemical literature these bonds are called hydrogen bonds.  This can be shown as: A - H...B;   where A most often represents oxygen or halogen, while B represents oxygen, halogen, or nitrogen.

The increase of the VWB energy is conditioned by electron-nuclear isomerization that proceeds, just as in the above cited cases, along the following scheme: A — H ... B ⇆ A-... (H ←B)+. In this case, the covalent bond and the VWB change into a polar VWB.

Just as in all the previously cited cases, the electronic-nuclear isomerization leads to the strengthening of the weak bonds and the weakening of the strong ones. As a matter of fact, the energy of the weak VWB (H ... B) increases, while the energy of the covalent bond A — H decreases. Thus, the increase of the VWB energy is more than 5 times greater than H ... B observed in the experiment, as compared with the common VWB, this is not paradoxical.

As a rule, the energy of a hydrogen bond (the increased VWB between H and B) is equal to10 - 150 kJ / mol.   (See table 6.6-1)

When this bond is formed, the distance between A and B, in the presence of hydrogen, is smaller (in spite of the hydrogen atom) than in analogous compounds without the hydrogen atom.

TABLE     6.6-1










     (F — H — F)-




     Cl — H ...Cl-




     Br — H...Br-




        I — H...I-










According to the theory, in compounds of the I3, XeF2, FHF, types, the bonding energy is defined by the energy of covalent bonds I2, XeF, FH. That is, the greater the energy of these bonds, the greater the bonding energy in compounds of the I3, XeF2 type.

In the row of compounds HF2, HCl2, HBr2, and HI2, the bonding energy decreases when transiting from left to right.. The covalent bonding energy decreases in row HF, HCl, HBr, and HI (See table  6.6-1).                                

Once more we mention the fact that all the atoms situated in molecules SF6, PCl5, XeF2, HF2, etc., cited above, the number of electrons in the outermost shells do not exceed 8; that is, the limitation of the number of covalent bonds that one atom can form with others is not violated. Likewise, other main precepts of the theory — the formation of chemical (covalent bonds, DABs, and VWBs — are not violated.

During the formation of one covalent bond, the number of electrons in the outermost electronic shell of the atoms being bonded increases by one unit. When DABs are formed, the number of electrons in the outermost shell of the donor atom does not change, while the number of electrons in the shell of the acceptor atom increase by 2 units. During VWB formation, the number of electrons in the outermost shells of the atoms being bonded does not change.

The main difference of the bonds, described in this section, is in the fact that the central atom in these compounds is bonded to other atoms with various types of bonds. For example — with a covalent bond and a common VWB, a covalent bond and a DAB, a single covalent bond and a double one, etc.; the weak bonds in these compounds strengthen while the strong bonds weaken. This is because in this case there is electronic isomerization, which proceeds reversibly. As a result of electronic isomerization, so-called electromers  are formed, i.e., electronic isomers where the bond used to be, for example, a covalent one, but then becomes a VWB.

With the decrease of the difference in the electronic energies of the electromers, the effect of energy decrease of the strong bond and the energy increase of the weak one is increased.

If the electronic energies of the electromers are equal, which is observed when the atoms are identical and bonded to the central atom, they have equalized energies and lengths. This is because, in this case, the transition speed in the reverse direction reaches the maximum — the greatest value for the given system.

Since the atoms' cores move 10 to 100 times slower than the electrons, they mainly occupy an intermediate position. Relatively, the number of electromers with weak bonds (like VWBs) abruptly decreases in the system, which decreases the speed of their eruption.

On the other hand, the high concentration of electromers where the nuclei occupy an intermediate position, there is an increase in nonproductive energy expenditure during thermal eruption.

Since the thermal stability of the bond (bonding energy) is defined experimentally by the total expenditure of the energy spent on the change of the electrons' potential and kinetic energies (where the potential energy decreases via the absolute value while the kinetic energy also decreases) and by the energy expenditure on the heating of the molecules that were not broken during the reaction (unproductive energy expenditure). With the increase of the unproductive expenditures, the experimentally defined bonding energy value increases.

The above-cited chemical bonds with common electronic isomerization that condition their peculiarities (first of all, the equalization of the length and the energy) and also their difference as compared with common bonds, allows giving these bonds the name of 'dynamic bonds'.

In the chemical bond section, these bonds look exotic, as if they are on the outside of the usual path of explaining the main chemical phenomenon — chemical bonding. However, when explaining the second main chemical phenomenon: chemical reactions, understanding the physical essence of dynamic bonds is the key. This fact once more underlines the good logic of their being singled out into a separate group with a separate name. Previously bonds, related to dynamic bonds, were cited separately in textbooks and scientific literature. 

Thus, hydrogen bonds were described as dipole bonds; bonds like SF6 and PF5,XeO3,XeF4 were explained as expanding the valence shell of the central atom that holds more than 8 electrons. The structures of compounds SO2, NO2, C6H6 were explained as resonance rules, as possible super-positioned structures written on the basis of the Lewis rules. Compounds like I3-, FHF-, Cl3-, BrCl2-, and some others were never discussed in textbooks.

In monographs (see Chemical Bonding Clarified Through Quantum Mechanics by G.C.Pimentel & R.D.Spratley) the bonds in I3-,Cl3 -,HF2 -,HCl2-,and some others, were explained in the framework of the theory of molecular orbitals as 3-electronic bonds. That is, various quantum-chemical suppositions were required in all the cited cases: super-positions, hybridization, molecular orbital theory, etc., which were regarded as unteachable.

The offered explanation is phenomenological and does not violate the theory of chemical bonding, which is also phenomenological in the framework of Introductory Chemistry.

According to the theory of chemical bonding, the outermost shells of atoms of electrons of the 2nd and 3rd periods in the table of elements, cannot contain more than 8 electrons.  

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 >>
Conclusions >>
Valence >>
Conclusions >>     
**Donor - acceptor Bond (DAB) >>   
Van der Waals Bond (VWB) >>    
Dynamic Bonds 
Conclusions >> 
Chapter 6 Textbook Questions