Comparación entre escenarios

ALCOHOLS ENERGY (kJ/mol) INORGANICS ENERGY (kJ/mol) CH3O-H...O(C2H5)2 10 (F — H — F)- 148 CH3O-H...N(C2H5)3 12 Cl — H ...Cl- ₅₆ Br — H...Br- ₅₂ PHENOLS I — H...I- ₄₈ C6H5O-H...O(C2H5)2 14.8 HO-H...OH2 20 C6H5O-H...N(C2H5)3 23.2 F-H...F-H 28

According to the theory, in compounds of the I3, XeF2and FHF, types, bonding energy is defined by the energy of covalent bonds I2, XeF and FH. That is, the greater the energy of these bonds, the greater the bonding energy in compounds of the I3and XeF2type.

In the row of compounds HF2, HCl2, HBr2and HI2, bonding energy decreases from left to right.. The covalent bonding energy decreases in the row of HF, HCl, HBr and HI (See table 6.6- 1).

We repeat that in all the atoms situated in molecules SF6, PCl5, XeF2, HF2, etc., as mentioned above, the number of electrons in the outermost shells does not exceed eight; 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, such as the formation of chemical bonds (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 increases by two units. During VWB formation, the number of electrons in the outermost shells of the atoms being bonded does not change.

The main difference between the compounds, as 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, Such as 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, placing electronic isomers where the bond used to be.

With the decrease of the difference in the electronic energies of the electromers, the effect of the 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 ten 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.

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 chemical bonds discussed above have common electronic isomerization that conditions their peculiarities (such as the equalization of the length and the energy) and differences from common bonds, which allows us to give these bonds the name 'dynamic bonds'.

In the chemical bond section, these bonds look exotic, as if they are outside the usual spectrum of explanation of the main chemical phenomenon, chemical bonding. However, when explaining the second main chemical phenomenon, chemical reactions, understanding the physical essence

of dynamic bonds is 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 discussed separately in textbooks and scientific literature.

Thus, hydrogen bonds were described as dipole bonds; bonds like SF6and PF5,XeO3,XeF4were explained as expanding the valence shell of the central atom that holds more than eight electrons. The structures of compounds SO2, NO2and C6H6were explained by resonance rules as possible super-positioned structures written on the basis of the Lewis rules. Compounds like I3-_{, FHF}-_, Cl3-_{and BrCl2}-_{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 similar others, were explained in the} framework of the theory of molecular orbitals as three-electron bonds. That is, various quantum- chemical suppositions were required in all the discussed 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 3}rd_{periods in the table of elements cannot contain more than eight electrons.}

Conclusions

The process of deepening the understanding of the physical nature of chemical bonding is currently continuing. Results so far are as follows. Bonds with which atoms are bonded into molecules can be divided into two types.

1) The first type of bond is the covalent bond. During its formation, the two outermost electrons (one from each of the atoms being bonded) rotate on a plane perpendicular to the axis connecting the bonding nuclei. These electrons will from now on be referred to as bonding electrons.

In the case of hydrogen, the nuclei being bonded are actually nuclei of hydrogen atoms. In all other cases, beginning with Li, when we speak of the atom's nucleus, we mean the nucleus and all the surrounding layers of the atom except the outermost shell.

When we speak of the effective charge of a nucleus we mean the charge that acts upon the bonding electron from the atom's nucleus and from all the other (nonbonding) electrons of the given atom.

There are two types of covalent bonds:

1a) If the effective charges of nuclei N1and N2are the same, such a bond is called a covalent homoatomic bond. The circle in which the bonding electrons rotate is situated at the same distance from the nuclei as the atoms being bonded. Such bonds are situated

in dual-atomic and multi-atomic molecules composed of similar atoms, like F2, Cl2, Na2, C6H14, etc.

1b) If the effective charges of nuclei N1and N2are different, then the circle in which the electrons rotate is closer to the atom with a greater nuclear charge. Such a bond is called a covalent heteroatomic bond. Such molecules as ClF, BrF and BrI are bonded with this kind of bond. If the effective nuclear charge of the atoms being bonded differs greatly, we are dealing with a super-polar or ionic bond. The atoms in salt molecules (NaCl, KF, LiF, etc.) are bonded with such bonds.

Even when forming super-polar bonds, electrons do not move from one atom to another. Moreover, the distance between the bonding electrons and the nuclei in the formed molecule (such as LiF) is smaller than the distance between the electrons and the nuclei in anions (such as Li-_{and F}-)_{; that is, during bond formation, the number of electrons in the outermost shells of} atoms, such as lithium (Li) and fluorine (F), increases by one electron.

The number of covalent bonds that one atom can form with other atoms is limited by the number of electrons situated in the outermost shell of the central atom. The number of bonds that atoms of groups I - IV of the table of elements can form is equal to the number of electrons in the outermost shells of these atoms. The number of covalent bonds that elements of groups V - VIII can form is limited by the maximal number of electrons that can be situated in the outermost electronic shells of atoms of this group (see the table of elements). When a covalent bond is formed, the number of electrons in the topmost shell of the atom increases by one.

2) The second type of bond is the donor-acceptor bond (DAB). Here, both bonding electrons belong to one of the atoms being bonded. The energy of this bond is about two times smaller than that of the covalent bond.

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 acceptor atom increases by two units.

The number of DABs that one atom can form with other atoms is limited, relative to the acceptor atom, by the number of electrons that the outermost shell of the given atom can contain. For atoms of periods 1, 2, 3 and 4 the maximal number of electrons in the outermost shells is equal to two, eight and eighteen respectively.

The number of DABs that the electrons' donor atoms can form is limited by the number of free electronic pairs (which do not take part in covalent bond formation).

Since energy gain is greater during covalent bond formation (because covalent bonds are stronger than DABs), atoms first form as many covalent bonds as possible.

The energy gain during chemical bond formation is conditioned by the approach of the electrons to the nuclei and by the increase of the effective charges of the nuclei of the atoms being bonded.

There is a greater energy gain in the case of a polar or super-polar bond than in the case of a covalent homo-polar bond, which is conditioned by a closer approach to the electrons of the nuclei and a greater effective charge (a greater initial FIE).

DABs, just like covalent bonds, can be polar. Just as in the case of covalent polar bonds, the strength of the polar DABs increases with the increase of the difference of the FIE of the atoms to be bonded.

That is, with the increase of the FIE difference, the possibility of DAB formation (stable compounds with DABs) increases.

3) Molecules are bonded between themselves via the third type of bond, the Van der Waals bond

(VWB), which is about ten times weaker than a covalent bond. During the formation of VWBs,

the number of electrons in the outermost shells of the atoms does not change.

The amount of energy that should be spent on breaking the bond (the strength of the bond) decreases according to the row thus: triple > double > heteroatomic > homoatomic covalent > DAB > VWB. The length of the bond increases in the same order.

The electrons of the outermost shells of the atoms take part in chemical bond formation. In the course of this formation, the potential and kinetic energies of the electrons change. The absolute value of the potential energies of the bonding electrons during bond formation increases, as do the kinetic energies of the bonding electrons. The energy gain (energy dispatch during bond formation) is conditioned by the kinetic energy increase (energy loss), which is two times smaller, relative to absolute value, than the potential energy increase. That is, the energy gain is equal to half of the potential energy gain.

Thus, bond formation is conditioned by the increase of the absolute value of the potential energies of the bonding electrons.

If various types of bonds bond an atom to similar atoms, such bonds become equal with respect to energy and length. The weak bonds become stronger while the strong ones become weaker. Analogously, the long bonds become shorter while the short ones become longer. The cause of this phenomenon is electronic isomerization.

In the course of isomerization, the electrons and the atoms' nuclei shift reversely; therefore, this type of bond can be singled out into a separate group of dynamic bonds. The strengthening of the weak bonds during isomerization explains the thermal stability of these compounds, since the thermal stability of a compound is defined by the energy of the weakest bond in the compound. The decrease of the strength of the thermal stability of the strong bond in the course of isomerization defines the key role of the dynamic bonds in chemical reactions.

The intermediate products of chemical reactions are compounds in which atoms of the reaction center are bonded with the different type of bonds with other atoms. First of all, the key role of

these bonds in chemical reactions is the basis for the singling out of this type of bond into a group called dynamic bonds.

The energy gain during molecule formation can be conditionally divided into two contributions: The first, smaller one is connected with the attraction of the nuclei to the bonding pair of electrons. The second, larger one is connected with the increase of the effective charge of the nuclei to be bonded during bond formation.

The amount of energy necessary for the thermal breaking of the chemical bond is about two times greater than the energy decrease during its formation out of atoms. This is because about half of the given energy is expended on heating the unbroken molecules.