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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.

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Valence Bond Theory

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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

  • 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.


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 energyfor 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 () bondcenters along the internuclear axis.  end-to-end overlap of orbitals

  • A pi () bondoccupies the space above and below the internuclear axis.  side-to-side overlap of orbitals


Figure 9.12. Sigma and Pi Bonding


Figure 9.10. An sp2 Hybridized C Atom


Figure 9.11. The s 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


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