Valence Bond Theory

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VALENCE BOND THEORY
The ‘mixing’ or ‘blending’ of atomic orbitals to accommodate the spatial requirements in a molecule is known as hybridization. Hybridization occurs to minimize electron pair repulsions when atoms are brought together to form molecules.

Possible hybridization schemes: 2nd row elements: sp sp2 sp3

3rd row elements also have:

dsp3

d2sp3

Each of these hybridzation schemes corresponds to one of the five fundamental VSEPR geometries.

Bonding arises from the overlap of orbitals. Sigma (σ) bonds arise from the ‘end-on’ overlap between adjacent orbitals. This leads to a region of high electron density along the inter-nuclear axis (cylindrically symmetrical).

Eg., 1s + 1s 2p + 2p

Pi (π) bonds arise from the ‘side-on’ overlap between adjacent orbitals. This leads to two regions of high electron density on opposite sides of the inter-nuclear axis (not cylindrically symmetrical).

Eg., 2p + 2p 3d + 3d
In all covalent bonding between atoms, there is one σ type bond. The remaining are π type.

Bonding involves the overlap of valence orbitals on the central atom with those of the surrounding atoms. Hybridization of pure atomic orbitals to form a special set of orbitals for use in bonding. Eg., CH4

Carbon:

2s22px12py1

one electron in each of four sp3

This hybridization allows carbon to form four rather than two covalent bonds and to orient them so as to minimize the e- pair repulsions (i.e., tetrahedrally).

Whenever a set of equivalent tetrahedral orbitals is required, the central atom adopts a set of four sp3 orbitals.

Bonding:

4 carbon-hydrogen σ bonds (sp3-1s)

Trigonal planar

sp2 hybridization

Whenever a central atom requires a trigonal planar geometry, it will adopt a set of three sp2 hybrid orbitals.

Eg., C2H4

Bonding:

4 carbon-hydrogen σ bonds (sp2-1s) 1 carbon-carbon σ bonds (sp2 – sp2) 1 carbon-carbon π bonds (2p – 2p)

Linear

sp hybridization

Whenever the central atom...
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