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Chemical Bond

Skills to develop

Chemical Bond

Chemical bond referrs to the forces holding atoms together to form molecules and solids. This force is of electric nature, and the attraction between electrons of one atom to the nucleus of another atom contribute to what is known as chemical bonds. Although electrons of one atom repell electrons of another, but the repulsion is relatively small. So is the repulsion between atomic nuclei.

Various theories regarding chemical bond have been proposed over the past 300 years, during which our interpretation of the world has also changed. Some old concepts such as Lewis dot structure and valency are still rather useful in our understanding of the chemical properties of atoms and molecules, and new concepts involving quantum mechanics of chemical bonding interpret modern observations very well.

While reading this page, you learn new concepts such as bondlength, bond energy, bond order, covalent bond, ionic bond, polar and non-polar bond etc. These concepts help you understand the material world at the molecular level.

Chemical bonds between identical atoms such as those in H2, N2, and O2 are called covalent bonds, in which the bonding electrons are shared. In ionic compounds, such as NaCl, the ions gather and arrange in a systematic fashion to form a solid. The arrangement of (blue) Na+ and (green) Cl- ions in a solid is shown in on the right here. The attraction force between ions are called ionic bonding. Matals such as sodium, copper, gold, iron etc. have special properties such as being good electric conductors. Electrons in these solids move freely throughout the entire solid, and the forces holding atoms together are called metallic bonds. To some extend, metals are ions submerged is electrons.

A Brief Past on Chemical Bond Concepts

Various concepts or theories have been proposed to explain the formation of compounds. In particular, chemical bonds were proposed to explain why and how one element reacted with another element.

In 1852, E. Frankland proposed the concept of valence. He suggested that each element formed compounds with definite amounts of other elements due to a valence connection. Each element has a definite number of valance.

Five years later, F.A. Kekule and others proposed a valence of 4 for carbon. Lines were used to represent valance, and this helped the development of organic chemistry. The structure of benzene was often quoted as an achievement in this development. More than 10 years later, J.H. van't Hoff and le Bel proposed the tetrahedral arrangement for the four valences around the carbon. These theory helped chemists to describe many organic compounds. In the mean time, chemical bonds were thaught to be electric nature. Since electrons have not been discovered as the negative charge carriers, they were thought to be involved in chemical bonds.

Following the discoveries of electrons by J.J. Thomson and R. A. Millikan, G.N. Lewis proposed to use dots to represent valence electrons. His dots made the valence electrons visible to chemists, and he has been credited as the originator of modern bonding theory. He published a book, in 1923, called Valence and the Structure of Atoms and Molecules.

X-ray diffractions by crystal allow us to calculate details of bondlength and bond angles. Using computers, we are able to generate images of molecules from the data provided by X-ray diffraction studies. These data prompted Linus Pauling to look at The Nature of Chemical Bond, a book that introduced many new concepts such as the resonance, electronegativity, ionic bond, and covalent bond.

In England, N.V. Sidgwick and H.E. Powell paid their attention to the lone pairs in a molecule. They developed the valence bond theory, the VSEPR (Valence Shell Electron Pair Repulsion) theory. An excellent summary of VSEPR is given on this link.

The application of quantum theory to chemical bonding gave birth to a molecular orbital theory.

In this and the few following modules, we will look at some of these concepts in detail.

Lewis Dot Structures

For the elements in the 2nd and 3rd periods, the number of valence electrons range from 1 to 8. Lewis dot structure for them are as indicated:
.   .    .    .     .     .    ..    ..
Li  Be  .B.  .C.   .N:   :O:  :F :  :Ne:  
    `         `     `     `    `     ``
Using dots, Lewis made the valence electron visible. The stability of noble gases is now associated with the 8 valence electrons around it. The stability of 8 valence electrons led him to conclude that all elements strive to acquire 8 electrons in the valence shell, and the chemical reaction takes place due to elements trying to get 8 electrons. This is the octet rule. For the hydrogen and helium atoms, 2 electrons instead of 8 are required.

For example, the octet rule applies to the following molecules:
H : H
(2 electrons)
. .
H : O : H
  . .
H : F :
. .
H : N : H
' '
. .
H : C : H
' '
: N ::: N : . .     . .
: O :: O :
. .    . .
: F : F :
. .           . .
: O :: C :: O :

To draw a Lewis dot structure, all the valence electrons are represented. A good way is to draw a type of dot for the valence electrons of one atom different from types in another. To do this on the computer screen using only type fonts is difficult, but you should draw a few by hand on paper.

When a dash is used to represent a bond, it represents a pair of electrons. Thus, in the following representations, a dash represents two electrons, bonding or lone pairs.

     _       _         _
  :S=O:     :O:H       :O:H
   |       _ |          | -
   O      :O:S:O:    :O:O:O:  
           " | "      " | "
            :O:H       :O:H
             "          "
These structures satisfy
the octet rule. Note
the two ways of drawing
the structures of

Exceptions to the Octet Rule

Elements in the 3rd and higher periods may have more than 8 valence electrons. A possible explanation for this is to say that these atoms have d-type atomic orbitals to accommodate more than 8 electrons. In the following molecules, the number of valence electrons in the central atoms are as indicated:
MoleculeSF6 PCl5 ICl3 XeF4
No. of valence
electrons for
central atom
12 10 10 12

Draw the Lewis dot structures for the above molecules, and count the number of valence electrons for the central atoms. For H2SO4, the S atom has 12 electrons in the structure shown on your right. Each dash represent a chemical bond, which has two electrons. There is a total of 6 bonds around the S atom, and therefore 12 electrons.

When B, Be, and some metals are the central atoms, they have less than 8 valence electrons. The following compounds do form, but the octet rule is not satisfied. These are electron defficient molecules.
MoleculeBeCl2BF3 BCl3 SnCl2
No. of valence
electrons for
central atom
46 6 6

Another case of exception to the octet rule are molecules with odd number of valence electrons. For example:
N O NO2 ClO2
No. of valence
electrons for
central atom

Isoelectronic Molecules and Ions

Counting the number of valence electrons often help us understand the formation of many molecules and ions. For example, all the following molecules have the 11 valence electrons:
NO CO- O2+ N2-
The charged molecule do exist under special circumstance.

The molecules of O2 are paramagnetic, and thus, they have unpaired electrons. The first dot structure does not agree with this observed fact, but the second one does. However, the second one does not obey the octet rule.
. .   . .
: O :: O :
.      .
: O ::: O :
No unpaired
octet rule

Later, you will learn that the molecular orbital (MO) theory provides a good explanation for the electronic configuration for O2.

Confidence Building Questions