Valence Bond Theory
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 in
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 in