SCH4U: and Orbital Hybridization

A) Valence Bond (VB) Theory (1927):  Developed by chemists Walter Heitler and Fritz London in 1927.  Proposed that a forms when two half-filled valence orbitals from two overlap.  The two overlapping orbitals form one “bonding orbital” = a volume of space between two atomic nuclei in which there is a high probability of finding the now spin paired bonding .

Figure 1: Ep and Covalent Bond Figure 2: Valence Bond +

Figure 3: Bonding in F2 – VB Theory

Figure 4: Bonding in H2O – VB Theory


VB Theory - Explaining ’s valence of 4:

Bonding in CH4 – VB Theory

A good bonding theory will explain all of the properties of the ’s structure including; , bond energy, bond angles and molecular shape.

Problem: for most , the VB theory does not match experimental observations.

Consider water: the bonds in H2O would involve the oxygen “2p” orbitals overlapping with the “1s” orbital. As a result, an expected bond angle of 90o. Experimental observations show that the bond angle in water is 104.5o.

Consider methane (CH4) the carbon would be using “2s” and “2p” orbitals to form bonds. The energy of “2p” orbital electrons is higher than the energy of “2s” orbital electrons, this should be reflected in the bond energies. In the real molecule all bond energies are equal. Also, this combination of orbitals can’t explain the tetrahedral shape of this molecule with bond angles of 109.5o.


B) Linus Pauling and “Orbital Hybridization” (1930):

Pauling proposed:  when bonding, the “pure” atomic orbitals of the central atom are replaced with new “hybrid orbitals”.  “s orbital” and “p orbitals” are reconfigured into “sp hybrid orbitals”

 hybridizing electrons from one “s” orbital and three “p” orbitals results in four identical hybrid orbitals called “sp3 hybrid orbitals” that are arranged in a tetrahedral pattern around the nucleus.  hybrid orbitals are intermediate in energy to the original “pure” orbitals.

C) Valence Bond Theory with Orbital Hybridization and Sigma (σ) and Pi (π) Bonds

3 In methane (CH4) the C atom uses “sp hybridization”. Each C-H bond is a formed when a C “2sp3” hybrid orbital overlaps with a H “1s” orbital.

Sigma Bonds (σ): A “sigma bond” is formed when two orbitals overlap in a direct “head-to-head” fashion. The “bonding orbital” formed by this type of overlap lies directly between and in the same plane as the two atomic nuclei. This very effective overlap results in a very strong bond.

The result is a methane molecule that matches the shape and bond angles that are both experimentally observed and predicted by VSEPR

Different hybridization schemes can be used to create different shapes/bonding patterns. -4-

Double Bond:  Two carbon atoms form a in ethene 2 (H2C=CH2). Here, the carbon atoms use “sp ” orbital hybridization.  “sp2” hybridization result from hybridizing one “s” orbital with two “p” orbitals producing three sp2 hybrid orbitals. This leaves one “p” orbital unhybridized.

 Three “sp2” orbitals arrange themselves in a planar triangle around nucleus. The unhybridized “p” is perpendicular to the “sp2” orbitals.

 One “sigma bond” forms when the two hybrid “2sp2” orbitals from each C overlap.  A “” forms when the two “2p” orbitals from each C overlap.

Pi Bonds (π): A “pi bond” is formed by the less effective “side-by-side” overlap of two “p” orbitals. This results in the “bonding orbital” forming above and below the plane of the nuclei. This less effective orbital overlap results in the pi bond being weaker than the sigma bond. Double bond = one sigma bond + one pi bond

Triple Bond:  Two carbon atoms for a in ethyne (HC≡CH). Here, the carbon atoms use “sp” orbital hybridization.

 “sp” hybridization results from hybridizing one “s” orbital with one “p” orbital to produce two “sp hybrid” orbitals. This leaves two “p” orbitals unhybridized.

 The two “sp” orbitals are arranged in a linear fashion around nucleus (180o apart). The two “p” orbitals are perpendicular to each other and the “sp” hybrids.

 A sigma bond forms when “2sp hybrid” orbitals from each C atom overlap. Two “pi bonds” form when the two unhybridized “2p” orbitals from each C each overlap.

Triple bond = one sigma bond + two pi bonds