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:
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
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.
4 carbon-hydrogen σ bonds (sp3-1s)
Whenever a central atom requires a trigonal planar geometry, it will adopt a set of three sp2 hybrid orbitals.
4 carbon-hydrogen σ bonds (sp2-1s) 1 carbon-carbon σ bonds (sp2 – sp2) 1 carbon-carbon π bonds (2p – 2p)
Whenever the central atom...
Please join StudyMode to read the full document