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


It has thus been proven experimentally that atoms are bonded into molecules; therefore, in order to break up a molecule into atoms, energy should be applied. For example, to break a hydrogen molecule (H2) into two atoms, it is necessary to heat the hydrogen to a temperature over 3,000 K or spend energy of more than 400 kJ/mol.

At 3,000 K the nuclei of hydrogen atoms have energy of about 20 kJ/mol.  As the atoms contain no other particles besides electrons, we can conclude that the energy spent on breaking the molecule into atoms, is actually used to change the electrons' energies.

The energy gain during bond formation is conditioned as follows: During molecule formation the electrons start moving in a field with a unified positive charge of the bonding nuclei. 

The electrons' energies are proportional to the square of the nuclear charges, and the inter-molecule repulsion energy is proportional to the charges in the first power. That is why, when the atoms approach each other at a certain point, the system's energy decreases.

When a molecule is formed out of atoms, the kinetic and potential energies of the electrons increase. The energy gain is conditioned by the fact that the increase of potential energy is twice greater than the increase of the electrons' kinetic energy.

That is, during the formation of a chemical bond, the electrons rotate around the greater positive charge and at a closer distance from the charge, which causes a greater stability of the system. This stability is defined by the energy that should be added to the system so as to break it up into constituent parts. In the case of molecules and atoms, it is the energy necessary to break up these systems into electrons and nuclei.

In a dual-atomic molecule, the positively charged nuclei are united with a circle of electrons rotating on a plane perpendicular to the axis connecting the nuclei. All the main parameters have been calculated for this system.  

As a result of calculations, it was found that:

1) When the bonding atoms have identical FIEs, the electrons' rotation plane is at the same distance from the bonding atoms. If the atoms' FIEs are different, the rotation plane will be shifted towards the atom's nucleus with a greater FIE. Two electrons partake in the formation of a chemical bond.

2) Both bonding electrons enter the outermost shells of the atoms being bonded.

3) During chemical bond formation, the number of electrons in  these outermost shells is increases by one.

4) During covalent bond formation, the bonding pair of electrons is composed of electrons from the atoms being bonded: one from each atom. The number of covalent bonds that one atom can form with other atoms  (atom's valence) is limited (due to inter-electronic repulsion) by the number of electrons that the given atom can connect to its outermost shell with energy gain.

5) The maximal number of electrons that the outermost shell can contain is equal to the number of electrons in the shells of the noble gas nearest to it in the table of elements. Noble gases cannot bond electrons in the outermost shell even when increasing the nuclear charge by 1 proton unit; therefore, they cannot form covalent bonds.

Thus, the number of covalent bonds, which an atom can form (atom's valence), is defined by the number of electrons in the outermost shell of the given atom (one outermost-shell electron is spent on the formation of one bond) and by the maximal number of electrons that can exist in the outermost shell of the given atom (amount of outermost-shell electrons is increased by 1 unit).

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 >>
**Donor - acceptor Bond (DAB) >>   
Van der Waals Bond (VWB) >>    
Dynamic Bonds >>
Conclusions >>