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Covalent Bonds



COVALENT BOND


  • For a covalent bond to form, two atoms must be located so that the orbital of one overlaps with the orbital of another; each orbital must contain a single electron.



  • When this happens, the two atomic orbitals merge to form two molecular orbitals. The orbital which is occupied by both the electrons, known as bonding MO. It results when orbitals with like sign overlap. The electrons are paired in a bonding MO i.e, they have opposite spins.


  • This arrangement is more stable than isolated atom; as a result formation of a bond is always accompanied by the evolution of energy.

This amount of energy per mole is called bond energy or bond dissociation energy when this amount of energy per mole is supplied to break the bond. (Bond dissociation energy = - Bond energy)

  • For a given pair of atoms, greater the overlap of atomic orbitals, stronger is the bond.

  • If AO’s of unlike signs overlap, an anti-bonding MO or MO* results which has a node (site of zero electron density) between the atoms and therefore has a higher energy than the individual AO’s.



σ AND π BONDS

  • Head-to-head overlap of AO’s give a sigma (σ) MO and the bond is called σ bond. The Head-to-head overlap of AO’s gives a sigma (σ) MO—the bonds are called σ bonds. The corresponding anti-bonding MO* is designated σ*. The imaginary line joining the nuclei of the bonding atoms is the bond axis, whose length is the bond length.


  • Two parallel p orbitals overlap side by side to form a pi (π) bond or a π * bond. The bond axis lies in a nodal plane (plane of zero electronic density) perpendicular to the cross-sectional plane of the σ bond. Single bonds are σ bonds. A double bond is one σ and one π bond. A triple bond is one σ and two π bonds a πx and a πy, if the triple bond is taken along the z-axis).

  • Although MO’s encompass the entire molecule, it is best to visualize most of them as being localized between pairs of bonding atoms. This description of bonding is called linear combination of atomic orbitals (LCAO).

  • Stabilities of molecules can be related to the bond order (B.O.), defined as


(Number of valence electrons in MO’s) – (Number of valence electrons in MO*’s)

B.O. =     __________________________________________________________________

2


  • The bond order is usually equal to the number of σ and π bonds between two atoms—in other words, 1 for a single bond, 2 for a double bond, 3 for a triple bond.



Difference between σ and π bonds

σ  bond

π bond

  1. Formed by head-to-head overlap of AO’s.

  1. Formed by lateral/sideways overlap of p /d orbitals

  1. Has cylindrical charge symmetry about bond axis.

  1. Has maximum charge density in the

cross-sectional plane of the orbitals

  1. Has free rotation.

  1. Does not have free rotation.

  1. Only one bond can exist between two atoms.

  1. One or two bonds can exist between two atoms.


Order of Energy

σ < π < π* < σ*

Hybridisation of atomic orbitals of carbon

AO’s of carbon hybridise in three ways:

  • In case of Alkanes like methane, four equivalent orbitals are formed by blending the 2s and the three 2p AO’s give four new hybrid orbitals, called sp3 HO’s.

  • In case of double bond or as it can be seen in alkenes, three equivalent orbitals are formed by blending the 2s and the two 2p AO’s to give three new hybrid orbitals, called sp2 HO’s and similarly for triple bonded carbons as in alkynes, two equivalent sp hybrid orbitals are formed by mixing of one 2s and one 2p orbital.

  • Only σ bonds and lone pair of electrons, not π bonds, determine hybridisation & molecular shapes.

  • Repulsion between pairs of electrons causes these HO’s to have the maximum bond angles and geometries. (For more details refer to chapter on Bonding)


TYPE OF HYBRIDISATION

BOND ANGLE

GEOMETRY

sp3

109.5º

Tetrahederal

sp2

120º

Triangular Planar

sp

180º

Linear


Tetrahederal geometry of methane


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