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

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


Valence is the number of electrons that an atom gains, loses, or shares when bonding with one or more other atoms. The valence in atoms is seen only when they are bonded.

One of the main questions in chemistry is:

Why is it that a strictly defined and limited number of atoms is bonded via covalent chemical bonding with one atom? 

Once more: In the course of a chemical experiment it was found that each atom is capable of bonding only a certain number of other atoms. Thus, for example, the atoms of elements of  group I in the table of elements are capable of connecting one hydrogen atom with the formation of compounds called metal hydrides. Atoms of the II, III, and IV groups can bond the number of hydrogen atoms that corresponds to the numbers of the group, i.e., 2, 3, and 4. The atoms of the V, VI, and VII groups bond only 3, 2, and 1 hydrogen atom respectively.

Another important question is:

How does our theory of chemical bonding explain the valence rules?

According to our theory, only the electrons in the outermost layer of the atom take part in chemical bond formation: (1) only one electron takes part in covalent bond formation. That is, the number of electrons on the outermost shell of each atom defines the possible number of bonds that the given atom can form. (2) In the course of chemical bond formation, the number of electrons in the outermost layer of each atom increases by one.       

Now let's cite an example on the basis of the second and third periods: How do our regularities explain the valence rules?

In row: Li, Be, B, C, N, O, F the number of electrons in the outermost layer of the atom increases by one unit when consecutively transiting from one member of the row to the next. Li has but one electron on the outermost shell. According to (1), Li can form one bond in its outermost shell; after bond formation, there will be 2 electrons. According to (2), Be, B, and C can form 2, 3, and 4 bonds respectively. And in their outermost shells, after bond formation, there will be 4, 6, and 8 electrons respectively. 

In the case of N, which has 5 electrons in the outermost layer, after the formation of 3 covalent dual-electronic bonds, its outermost layer will contain 8 electrons (5 initial ones and 3 additional ones).

As indicated in the previous section, atoms, with 8 electrons in the outermost shells, do not form covalent bonds; that is, out of 5 electrons of the N-atom in the outermost layer for bond formation, the N-atom is capable of utilizing only three.

A homopolar covalent bond is one whose dipole moment has a value close to zero. It can be either homoatomic (Cl2, Na2) or heteroatomic (LiNa). A heteropolar covalent bond is one whose dipole moment is higher than zero (NaCl, BaCl2, etc.).

A molecule of NaCl is formed of atoms of sodium and chlorine. Chlorine has 8 electrons in its outermost shell, while sodium has only 2. That is, the outermost shell of the sodium is unfilled. On the other hand, only 2 electrons from the filled chlorine shell take part in the formation of the heteropolar bond Na - Cl.   Six electrons (3 pairs) in the outermost shell of Cl do not take part in bond formation. That is, in the case of molecule NaCl, the sodium can bond 6 more electrons while the chlorine can offer 6 electrons for bonding. Thus, one NaCl molecule can connect another NaCl molecule with the formation of a dimer Na2Cl2, which can be illustrated as a molecule:                          

Na : Cl
** **
Cl : Na

In this molecule, the electrons, indicated via two dots (:) belonged separately to sodium and chlorine atoms before the bonding. The electrons indicated with dots (*) belonged to chlorine in molecule NaCl before the bonding.

The bond between two atoms, which is formed at the expense of the electrons of one atom are called donor-acceptor bonds (DAB). Here the bonding pair of electrons also enters the outermost shells of both atoms being bonded. The energy gain, just as in covalent bonding, is attained at the expense of the concentration of the positive charge. Its difference lies in the fact that during bond formation, not one but both electrons, which belonged to one atom (a donor) before the formation, enter the shell of another atom (an acceptor).

The Na2Cl2 molecule was received experimentally, so was the reaction energy 2NaCl → Na2Cl2. This energy is equal to 186 kJ/mol (one DAB being equal to 93 kJ/mol). The bonding energy of a NaCl monomer is equal to 410 kJ/mol, while the bonding energies of Na - Na   and Cl - Cl comprise, relatively, 75 kJ/mol and 238 kJ/mol, i.e., the DAB is much weaker than the covalent bond.

One of the main reasons for the comparatively weak DAB is as follows. When calculating the bonding energy, the initial atoms' electronic energy, taking part in bond formation, is subtracted from the molecule's calculated energy. The ionization energy of the two non-bonding chlorine electrons in NaCl exceed the sum of the FIEs of both sodium and chlorine atoms.

That is, the energy gain (difference between the electronic energies of the divided and bonded atoms via chemical bonding) in the case of the DAB is smaller at the expense of the energy increase of the divided atoms. The DAB is a well-known kind of chemical bonding. It is due to the DABs that saturated molecules with non-binding (lone pair) electrons are connected to atoms in molecules with unfilled outermost shells. The above-cited example with NaCl is an example of such a bond.

Historically, molecules containing electron donor atoms are called ligands and, in the capacity of acceptors, are the elements of the 4th, 5th, and 6th periods in the table of elements. The saturation of the outermost electronic shells of these elements is possible when there are more electrons than in the outermost shell of the elements of the 2nd and 3rd periods.

Thus, for example, the saturation of the outermost shells with nickel (Ni), cobalt (Co), rhodium (Rh), is possible when each of these have 18 electrons. The atoms of these metals have only 10 electrons in the outermost electronic shells, but can bond 4 molecules of carbon monoxide via the DAB when in the atomic state.

Each molecule of carbon monoxide increases the number of electrons in the outermost shell of a metal by 2. The outermost shell of the metal molecule M(CO)4 contains 18 electrons.

Historically, such compounds have been called coordinate compounds. Of these compounds we single out the M (coordinate metal) and molecules (in this case —CO), or ligands.

Now let's see what kind of interaction we can expect between atoms whose outermost electronic shells are filled. In this case, no DAB is formed, however, there is a slight energy gain at the expense of the positive charge concentration around which the electrons rotate. Bonds caused by this energy gain have been called Van der Waals bonds (VWB). The energy of these bonds is about ten times smaller than that of the covalent bonds. Such bonds are most common between molecules where the outermost shells of the atoms are completely filled with electrons. For example, bonds between halogen molecules:

F 2 ... F2           Cl2 ... Cl2


The VWBs are connected among themselves and to the noble gas atoms:

Ne             Ne

The filling of the outermost shell with atoms connected with covalent bonds can occur not only at the expense of the connection of new atoms, but also at the expense of the formation of multiple bonds. Thus, for example, in the case of carbon (C), besides the existence of compound:                        

                                                                                                H      H

                                                                                          H : C  -  C : H

                                                                                                H      H

where the shells of  both carbon atoms are filled at the expense of  the hydrogen atoms' electrons, there are also compounds of this type: 

H: C :: С : Н          and           H : C : : : C : H

In these compounds the outermost electronic carbon shells are filled partially by electrons at the expense of the electrons of each of the connecting carbon atoms. In ethylene the carbons are connected between themselves with a double bond, and in acetylene - with a triple one.    Unlike cases with single bonds, in this case there is repulsion between the bonding pairs of electrons; and we should expect a decrease in the energy gain at bond formation.

In other words, if we compare the energy gain of six bonds formed between carbon and hydrogen in a carbon-carbon bond in ethane, with the energy gain at the expense of the formation of four carbon-hydrogen bonds and two carbon-carbon bonds in ethylene, the latter will be greater.

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