%D, %d %M %y
Time: %h~:~%m

Home  / GENERAL CHEMISTRY Textbook / Chapter 6. Molecule structure / Conclusions

Conclusions

At present the process of deepening the understanding of the physical nature of chemical bonding is continuing. Results, already gained in this respect, 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 in a plane perpendicular to the axis connecting the bonding nuclei. These electrons will further 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 excepting the outermost shell.

When we speak of the effective charge of a nucleus we mean the charge, which 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 N1 and N2 are 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 N1 and N2 are 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, 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 the super-polar bond, the electrons do not transit from one atom to another. Moreover, the distance between the bonding electrons and the nuclei in the formed molecule (for example LiF) is smaller than the distance between the electrons and the nuclei in anions 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 that are 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 forming one covalent bond, the number of electrons in the topmost shell of the atom increases by one electron.

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 twice smaller than that of the covalent bond.

When forming DABs 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 2 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 2, 8.8, and 18 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 the energy gain is greater during covalent bond formation (covalent bonds are stronger than DABs), first the atom forms all the possible covalent bonds.

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 that 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 as covalent bonds, can be polar. Just like 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 and the kinetic energies of the bonding electrons increases. 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 in 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 (move) 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 (bigger 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 used rather uselessly on heating the unbroken molecules.