Valence Bond Theory
How do bonds form?
• The valence bond model or atomic orbital model
was developed by Linus Pauling in order to
explain how atoms come together and form
molecules.
• The model theorizes that a covalent bond forms
when two orbitals overlap to produce a new
combined orbital containing two electrons of
opposite spin.
• This overlapping results in a decrease in the
energy of the atoms forming the bond.
• The shared electron pair is most likely to be found
in the space between the two nuclei of the atoms
forming the bonds.
Example H2
1s
1s
H
H
1s
Overlapping of the
1s orbitals
Covalent Bond
H-H
• The newly combined orbital will contain an electron
pair with opposite spin just like a filled atomic orbital.
Example HF
• In hydrogen fluoride the 1s orbital of the H will overlap
with the half-filled 2p orbital of the F forming a covalent
bond.
1s
2p
+
+
H
F
Overlapping of the1s and 2p orbitals
+
Covalent Bond
H-F
Other Points on the Valence Bond Theory
• This theory can also be applied to molecules
with more than two atoms such as water.
• Each covalent bond results in a new
combined orbital with two oppositely
spinning electrons.
• In order for atoms to bond according to the
valence bond model, the orbitals must have
an unpaired electron.
Covalent Bonding: Orbitals
Hybridization
• The mixing of atomic orbitals to form
special orbitals for bonding.
• The atoms are responding as needed to
give the minimum energy for the
molecule.
sp3 Hybridization
The experimentally known structure of CH4 molecule
can be explained if we assume that the carbon atom
adopts a special set of atomic orbitals. These new orbital
are obtained by combining the 2s and the three 2p
orbitals of the carbon atom to produce four identically
shaped orbital that are oriented toward the corners of a
tetrahedron and are used to bond to the hydrogen atoms.
Whenever a set of equivalent tetrahedral atomic orbitals
is required by an atom, this model assumes that the atom
adopts a set of sp3 orbitals; the atom becomes sp3
hybridized.
Figure 9.5. An Energy-Level Diagram Showing the
Formation of Four sp3 Orbitals
Figure 9.2. The Valence Orbitals on a Free Carbon
Atom: 2s, 2px, 2py, and 2pz
Figure 9.3. The Formation of sp3 Hybrid Orbitals
Figure 9.6. Tetrahedral Set of Four sp3 Orbitals
Figure 9.7. The Nitrogen Atom in Ammonia is sp3 Hybridized
Figure 9.9. An Orbital Energy-Level Diagram for sp2 Hybridization
Figure 9.8. The Hybridization of the s, px, and py Atomic Orbitals
• A sigma () bond centers along the
internuclear axis.  end-to-end overlap
of orbitals
• A pi () bond occupies the space above
and below the internuclear axis.  sideto-side overlap of orbitals
H

H
C C
H
H
Figure 9.12. Sigma and Pi Bonding
Figure 9.10. An sp2 Hybridized C Atom
Figure 9.11. The  Bonds in Ethylene
Figure 9.13. The Orbitals for C2H4
Figure 9.16. The Orbital Energy-Level Diagram for the
Formation of sp Hybrid Orbitals on Carbon
Figure 9.14. When One s Orbital and One p Orbital are
Hybridized, a Set of Two sp Orbitals Oriented at 180
Degrees Results
Figure 9.17. The Orbitals of an sp Hybridized Carbon Atom
Figure 9.18. The Orbital Arrangement
for an sp2 Hybridized Oxygen Atom
Figure 9.15. The Hybrid Orbitals in the CO2 Molecule
Figure 9.19. The Orbitals for CO2
Figure 9.20. The Orbitals for N2
Figure 9.21. A Set of dsp3 Hybrid Orbitals on a Phosphorus Atom
Figure 9.23. An Octahedral Set of d2sp3 Orbitals on a Sulfur Atom
Figure 9.24. The Relationship of the Number of Effective Pairs,
Their Spatial Arrangement, and the Hybrid Orbital Set Required
Figure 9.46. A Benzene Ring
Figure 9.47. The Sigma System for Benzene
Figure 9.48. The Pi System for Benzene
The Localized Electron Model

Three Steps:
Draw the Lewis structure(s)

Determine the arrangement of electron
pairs (VSEPR model).

Specify the necessary hybrid orbitals.
Figure 9.45. The Resonance Structures for O3 and NO3-
Paramagnetism

unpaired electrons

attracted to induced magnetic field

much stronger than diamagnetism