Kossel-Lewis Approach to Chemical Bonding
Table of contents
- Introduction to Kossel-Lewis Approach to Chemical Bonding
- Octet Rule
- Condition for writing the Lewis dot structures of molecules
- Types of Bonds
- Examples of Covalent Bond
- Writing Lewis structures
In 1916 Kossel and Lewis succeeded in giving successful explanation based upon the concept of electronic configuration of noble gases about why atoms combine to form molecules. Atoms of noble gases have little or no tendency to combine with each other or with atoms of other elements. This means that these atoms must be having stable electronic configurations. The electronic configurations of noble gases are given in figure below.
Due to the stable configuration, the noble gas atoms neither have any tendency to gain nor lose electrons and, therefore, their combining capacity or valency is zero. They are so inert that they even do not form diatomic molecules and exist as monoatomic gaseous atoms.
All atoms other than noble gases have less than eight electrons in their outermost shells. In other words, the outermost shells of these atoms do not have stable configurations. Therefore, they combine with each other or with other atoms to achieve stable noble gas electronic configurations. These elements undergo electronic re-arrangements to attain stable noble gas configurations. Therefore
“The tendency or urge of atoms of various elements to attain stable configuration of eight electrons in their valence shells, is the cause of Chemical combination” and
“The principle of attaining maximum of eight electrons in the valence shell of atoms, is called octet rule.”
Lewis introduced simple symbols to denote the electrons present in the outer orbit of atom, these electrons are known as valence electrons. These symbols are known as Electron Dot Symbols and the structure of compound is known as Lewis Dot Structure.
(Note: Number of dots around the symbol is equal to number of electrons.)
Lewis Structures and Covalent Bond condition for writing the Lewis dot structures of molecules
Conditions for writing the electron dot (or Lewis) structures of covalent molecules are:
- Sharing of an electron pair between the atoms results in the bond formation.
- At the time of bond formation, each bond consists of two electrons which are contributed by each one of the combining atoms.
- When the combining atoms forms bond, because of sharing of electrons they achieve the stable outer shell noble gas configurations. In other words, octets of both the atoms get completed.
Electron dot (or Lewis) structures of covalent molecules are written in accordance with octet rule. According to this, all the atoms in the formula will have a total of eight electrons in their valence shell except the Hydrogen atom. Hydrogen will have only two electrons because only two electrons complete its first shell as in helium. Thus the elements of group 17 (containing seven valence electrons) such as Cl would share one electron to attain stable octet; the elements of group 16 (containing six valence electron) such as O and S would share two electrons; the elements of group 15 (containing five valence electrons) would share three electrons and so on.
For Example, oxygen (with six electrons in the valence shell) completes its octet by sharing its two electrons with two hydrogen atoms to form a water molecule as shown below.
Similarly, nitrogen has five electrons in its valence shell and shares with three hydrogen atoms to form ammonia.
Similarly, carbon has four electrons in its valence shell and shares with four chlorine atoms to form carbon tetrachloride (CCl4) molecule as shown below.
Lewis structure for molecules having multiple covalent bonds. If the normal valence of an atom is not satisfied by sharing single electron pair between atoms, the atoms may share more than one electron pair between them.
- Single covalent bond forms when two atoms share one electron pair and is represented by one dash (–)
- Double covalent bond when two atoms share two electron pairs and is represented by two dashes ( = ).
- Triple covalent bond forms when two atoms share three electron pairs and is represented by three dashes ( ≡ ).
Let us discuss some examples of molecules having double and triple bonds.
In the formation of oxygen molecule, each oxygen atom has six electrons in the valence shell and requires two electrons to complete their octet. Therefore the atoms contribute two electrons each for sharing to form oxygen molecule, two electron pair are shared and hence there is a double bond between the two oxygen atoms.
Carbon has four valence electrons and oxygen has six. To complete the octets, carbon shares two of its valence electrons with one oxygen atom and two with other oxygen atom.
Thus, there are two double bonds in CO2 molecule.
In the formation of a nitrogen molecule each of the two nitrogen atoms having five valence electrons, provides three electrons to form three electron pairs for sharing. Thus, a triple bond is formed between the two nitrogen atoms.
In ethylene, each carbon atom shares two of its valence electron with two hydrogen atom and remaining two electrons with the second carbon atoms. So there is a double bond between the carbon atoms.
The following steps are adopted for writing the Lewis dot structures or Lewis structures:
Step 1: First, we calculate the required number of electrons for drawing the structure by adding the valence electrons of the combining atoms. For Example, in methane, CH4 molecule, there are 8 valence electrons (in which 4 belongs to carbon while other 4 to H atoms).
Step 2: Each negative charge i.e. for anions, we add an electron to the valence electrons and for each positive charge that is, for cations we subtract one electron from the valence electrons.
Step 3: Using the chemical symbols of the combining atoms and constructing skeletal structure of the compound, divide the total number of electrons as bonding shared pairs between the atoms in proportion to the total bonds.
Step 4: The central position in the molecule is occupied by the least electronegative atom occupies Hydrogen and fluorine generally occupy the terminal positions.
Step 5: After distributing the shared pairs of electrons for single bonds, the remaining electron pairs are used either for multiple bonds or they constitute lone pairs.
The basic requirement is that each bonded atom gets an octet of electrons.
Example No. 1 Lewis formula for carbon monoxide, CO
Step 1: Counting the total number of valence electrons of carbon and oxygen atoms:
C (2s22p2) + O (2s22p4)
4 + 6 = 10 that is, 4(C) + 6(O) = 10
Step 2: The skeletal structure of carbon monoxide is written as: CO
Step 3:.Drawing a single bond between C and O and completing octet on O, the remaining two electrons are lone pair on C.
Step 4: This does not complete the octet of carbon, and hence we have triple bond between C and O atom
This satisfies octet rule for both the atoms.
Example No. 2 Lewis Structure of nitrite, No2-
Step 1: Counting the total number of valence electrons of one nitrogen atom, two oxygen atoms and the additional one negative charge (equal to one electron).
Total Number of valence electrons are:
N (2s22p3) + 2O (2s22p4) + 1 (negative charge)
5 + 2 6 + 1 = 18
Step 2: The skeletal structure of nitrite ion is written as
O N O
Step 3: Drawing a single bond between nitrogen and each oxygen atom
Step 4: Complete the octets of atoms.
This structure does not complete octet on N if the remaining two electron constitute of a lone pair on it. Therefore, we have double bond between one N and one of the two O atoms.
The Lewis structure is