Notes On Organic Compounds: Shapes - CBSE Class 11 Chemistry
The tetravalence of carbon can  be explained from its excited valence shell electronic configuration. The outermost shell of an atom is known as the valence shell. Electrons present in the valence shell (valence electrons) take part in the chemical reaction.



The electronic configuration of carbon is one 1s2 2s2 2px12py1.

All elements tend to complete octet in their valence shell to become stable. Since carbon has four electrons in its valence shell to complete its octet, it forms four covalent bonds is called as tetravalence of carbon.

The formation of molecules and their shapes can be explained by considering the hybridisation of s and p orbitals respectively. In ethene each carbon atom undergoes sp2 hybridization. 2sp2 hybrid orbitals of each carbon atom overlaps with each of the 1s orbitals of two hydrogen atoms  to form two sigma bonds respectively.

The remaining sp2 hybrid orbital of a carbon atom overlaps co-axially with a similar sp2 hybrid orbital of the other carbon atom to form a strong carbon-carbon sigma bond by sp²-sp² overlap. The unhybridised (or) pure 2pz-orbitals of both the carbon atoms overlap laterally to form a weak carbon-carbon pi-bond. Pi bond doesn’t account for the shape of the molecule.

All the six atoms containing the sigma bonds lie in one plane and the pi bond lies perpendicular to the plane. Thus the shape of the molecule is trigonal planar.



Hybridization influences bond length and bond strength in organic compounds. Greater is the s-character, closer is the orbital to the nucleus and hence higher is the electronegativity of that carbon.

Bond strength of sp-hybrid orbital is greater than sp2 and sp3 hybrid orbitals. Thus, the greater the s-character, the greater is the electronegativity. sp is more electronegative followed by sp2 and sp3.

sp > sp2 > sp3

The bonds formed by lateral overlapping of pure p-orbitals are known as pi bonds. The pure  p-orbitals of both the carbon atoms which  are mutually parallel to each other  form  a pi-bond that is perpendicular to the plane of the molecule.

Rotation of C-C around the double bond  is  strongly restricted, as the pi bonds prevent the free rotation  around the carbon -carbon bond.



In a pi bond the electron charge cloud is present above and below the plane of the bonding atoms. Hence, these electrons are more exposed and susceptible to attack by electron seeking reagents.
In general, pi bonds provide the most reactive centres in unsaturated molecules.

For example, the pi  electrons are more tightly held by the carbon atoms in acetylene as the carbon atoms are sp hybridised compared to the pi electrons of ethylene in which the carbon atoms are sp2 hybridised. Hence acetylene is less reactive than ethylene.

Summary

The tetravalence of carbon can  be explained from its excited valence shell electronic configuration. The outermost shell of an atom is known as the valence shell. Electrons present in the valence shell (valence electrons) take part in the chemical reaction.



The electronic configuration of carbon is one 1s2 2s2 2px12py1.

All elements tend to complete octet in their valence shell to become stable. Since carbon has four electrons in its valence shell to complete its octet, it forms four covalent bonds is called as tetravalence of carbon.

The formation of molecules and their shapes can be explained by considering the hybridisation of s and p orbitals respectively. In ethene each carbon atom undergoes sp2 hybridization. 2sp2 hybrid orbitals of each carbon atom overlaps with each of the 1s orbitals of two hydrogen atoms  to form two sigma bonds respectively.

The remaining sp2 hybrid orbital of a carbon atom overlaps co-axially with a similar sp2 hybrid orbital of the other carbon atom to form a strong carbon-carbon sigma bond by sp²-sp² overlap. The unhybridised (or) pure 2pz-orbitals of both the carbon atoms overlap laterally to form a weak carbon-carbon pi-bond. Pi bond doesn’t account for the shape of the molecule.

All the six atoms containing the sigma bonds lie in one plane and the pi bond lies perpendicular to the plane. Thus the shape of the molecule is trigonal planar.



Hybridization influences bond length and bond strength in organic compounds. Greater is the s-character, closer is the orbital to the nucleus and hence higher is the electronegativity of that carbon.

Bond strength of sp-hybrid orbital is greater than sp2 and sp3 hybrid orbitals. Thus, the greater the s-character, the greater is the electronegativity. sp is more electronegative followed by sp2 and sp3.

sp > sp2 > sp3

The bonds formed by lateral overlapping of pure p-orbitals are known as pi bonds. The pure  p-orbitals of both the carbon atoms which  are mutually parallel to each other  form  a pi-bond that is perpendicular to the plane of the molecule.

Rotation of C-C around the double bond  is  strongly restricted, as the pi bonds prevent the free rotation  around the carbon -carbon bond.



In a pi bond the electron charge cloud is present above and below the plane of the bonding atoms. Hence, these electrons are more exposed and susceptible to attack by electron seeking reagents.
In general, pi bonds provide the most reactive centres in unsaturated molecules.

For example, the pi  electrons are more tightly held by the carbon atoms in acetylene as the carbon atoms are sp hybridised compared to the pi electrons of ethylene in which the carbon atoms are sp2 hybridised. Hence acetylene is less reactive than ethylene.

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