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Electron Structure of Carbon

Carbon is an element of second period of periodic table. Its atomic number is 6 and mass number is 12. It is represented as 12 6C.

The ground state electronic configuration of carbon is:

6C (Ground state): 1s2 2s2 2px 1 2pz 0

Carbon atom has 4 electrons in its valence shell and, therefore, it can attain a noble gas configuration either by losing or gaining or sharing 4 electrons. But the loss or gain of 4 electrons by the carbon atom to form highly charged C+4 or C-4 ions would require a very large amount of energy which is not ordinarily available during a chemical reaction. Therefore, carbon is unable to form ionic bonds and as such it can participate only in the formation of covalent bonds.

The tendency of carbon atom to form covalent bonds only is also justified on the basis of its electronegativity value, which is 2.5. Since the carbon atom lies in group I of the periodic table, its electronegativity is neither very low nor very high. The electronegativities of other elements such as H (2.1), 0 (3.5), N (3.0), Cl (3.0) and S (2.5) which are generally linked with carbon in organic compounds are not very much different from that of carbon. Because of this small electronegativity difference between carbon and the other elements bonded to it, the bonds formed are predominantly covalent.

The electronic configuration of carbon suggests that it should be bivalent i.e,, should show a valency of two because of the presence of two half-filled 2p orbitals (i.e., 2px and 2py orbitals) in its valence shell. All the known compounds of carbon confirm that carbon is tetracovalent. In order to account for tetravalency, it is believed that during the process of bond formation, which is energy releasing process, the two electrons in the 2s orbital get unpaired and out of them one is promoted to empty 2p z orbital. This is called excited state. The electronic configuration of carbon in the excited state is :

6C (excited state): 1s2 2s2 2px 1 2pz 1

Carbon atom in the excited stale has four half-filled orbitals (2s1 1p/ 2p 12p/). Since all these orbitals would be available for the foT71UJtion of covalent bonds with other atoms, carbon would exhibit a covalency of four.


It was proposed that the electrons involved in the process of combination are outer shell electrons, therefore, these are called valence electrons. G.N. Lewis introduced the simple notations to represent the valence electrons in an atom. These notations are called Lewis symbols or electron dot symbols. According to these notations,

(i) the symbol of the dement represents the nucleus as well as the electrons in the inner shells.

(ii) the electrons in the outer shell are represented by the dots surrounding the symbol.

For example, the Lewis symbols of carbon shall be written as:

The electrons of valence shell are represented by dots (•) or by cross (X). The Lewis symbols for the other elements of 2nd period are:

Significance of Lewis Symbols

The number of dots in the Lewis symbol represents the number of valence electrons. The common valencies of the elements can be calculated from the valence electrons. The common valency of the element is either equal to the number of dots or 8 minus the number of dots. For example, the common valencies of Li, Be, B and C are 1, 2, 3 and 4 respectively while those of N, 0 , F and Ne are 8 minus number of dots, i.e., 3, 2, 1 and 0 respectively.

Lewis Representation of Simple Molecules and Polyatomic Ions

The Lewis dot structures provide a picture of bonding in simple molecules and polyatomic ions in terms of shared electron pairs. The following basic steps are generally used for writing Lewis dot structures.

(i) Add the valence electrons of the combining atoms. This gives the total number of electrons for writing the structure.

(ii) For polyatomic anions, add one electron to the total number for each unit negative charge. For example, in CO3 2- ion, the total number of electrons are calculated as

Total electrons = (4e- from carbon)

+ 3 x (6e- from each 0 atom)

+ 2 (1e- for each negative charge)

= 4 + 3 x 6 + 2 x 1 = 24 electrons.

(iii) For polyatomic cation subtract one electron for each unit positive charge.

(iv) Write the skeleton structure by placing the least electronegative atom in the centre and more electronegative atoms on the terminal positions. While doing this, some intelligent guess work is also required.

(v) Now distribute the electrons properly as shared pairs in proportion to the total bonds.

(vi) After accounting for the shared pair of electrons for single bonds, the remaining electrons are utilized either for multiple bonding or for indicating the lone pairs.

Illustrations Let us write the Lewis dot structure for CO and HCN molecules


1. Carbon monoxide (CO) molecule

Step I. Total number of valence electrons = ( 4e- from C)+ (6e- from 0) = 10e-

Step II. The skeletal structure is CO

Step III. Draw single bond between carbon and oxygen atom (one shared pair) and complete the octet on 0 atom. The remaining two electrons are the lone pair on C atom.

Step IV. This does not complete the octet on carbon atom. Hence we have to resort to multiple bonding between carbon and oxygen to satisfy the octet. Hence, the Lewis formula for carbon monoxide is

2. Hydrogen cyanide (HCN) molecule

Step I. Total number of valence electrons

= (le- from H)+ (4e- from C)+ (5e- from N)

= 10e-

Step II. The skeletal structure is H C N

Step III. Put one shared pair of electrons between H and C and one shared pair of electrons between C and N. The remaining electrons are lone pairs on C and N atoms.

Step IV. Since octet of C and N are not complete, thus multiple bonding is required between C and N atoms. For completing their octets triple bonding is required between C and N atoms. Thus, the structure is

Let us now, write the Lewis structures of some more molecules and polyatomic ions.