Valence Bond Theory

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Valence Bond Theory
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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.

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Example H2

• The newly combined orbital will contain an
  electron pair with opposite spin just like a
  filled atomic orbital.

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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.

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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.

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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.

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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.

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Figure 9.5. An Energy-Level Diagram Showing the
Formation of Four sp3 Orbitals
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Figure 9.2. The Valence Orbitals on a Free Carbon
Atom 2s, 2px, 2py, and 2pz
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Figure 9.3. The Formation of sp3 Hybrid Orbitals
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Figure 9.6. Tetrahedral Set of Four sp3 Orbitals
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Figure 9.7. The Nitrogen Atom in Ammonia is sp3
Hybridized
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Figure 9.9. An Orbital Energy-Level Diagram for
sp2 Hybridization
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Figure 9.8. The Hybridization of the s, px, and
py Atomic Orbitals
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• 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. ? side-to-side overlap of
  orbitals

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Figure 9.12. Sigma and Pi Bonding
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Figure 9.10. An sp2 Hybridized C Atom
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Figure 9.11. The s Bonds in Ethylene
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Figure 9.13. The Orbitals for C2H4
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Figure 9.16. The Orbital Energy-Level Diagram for
the Formation of sp Hybrid Orbitals on Carbon
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Figure 9.14. When One s Orbital and One p Orbital
are Hybridized, a Set of Two sp Orbitals Oriented
at 180 Degrees Results
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Figure 9.17. The Orbitals of an sp Hybridized
Carbon Atom
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Figure 9.18. The Orbital Arrangement for an sp2
Hybridized Oxygen Atom
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Figure 9.15. The Hybrid Orbitals in the CO2
Molecule
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Figure 9.19. The Orbitals for CO2
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Figure 9.20. The Orbitals for N2
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Figure 9.21. A Set of dsp3 Hybrid Orbitals on a
Phosphorus Atom
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Figure 9.23. An Octahedral Set of d2sp3 Orbitals
on a Sulfur Atom
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Figure 9.24. The Relationship of the Number of
Effective Pairs, Their Spatial Arrangement, and
the Hybrid Orbital Set Required
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Figure 9.46. A Benzene Ring
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Figure 9.47. The Sigma System for Benzene
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Figure 9.48. The Pi System for Benzene
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The Localized Electron Model

• Three Steps
• Draw the Lewis structure(s)
• Determine the arrangement of electron pairs
  (VSEPR model).
• Specify the necessary hybrid orbitals.

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Figure 9.45. The Resonance Structures for O3 and
NO3-
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Paramagnetism

• unpaired electrons
• attracted to induced magnetic field
• much stronger than diamagnetism