1. COVALENT BONDING: ORBITALS
Chapter 9

2. Hybridization
The mixing of atomic orbitals to form special molecular orbitals for bonding.
The atoms are responding as needed to give the minimum energy for the molecule.

3. Molecular Geometry & Hybridization
Parent geometry determines the hybridization.
Molecular structure is the actual geometry.

4. Energy-level diagram showing the formation of four sp3
hybrid orbitals.

5. One 2s and three 2p orbitals hybridize to form a
new set of sp3 hybrid orbitals.

6. sp3 hybrid orbital 4 effective electron pairs. tetrahedral geometry.
109.5 o bond angle.

7. The tetrahedral set of four sp3 orbitals of the
carbon atom share one electron each with the four
hydrogen atoms to make a methane molecule.

8. The nitrogen atom in ammonia is sp3 hybridized.

9. A sigma () bond centers along the internuclear axis.
A pi () bond occupies the space above and below the internuclear axis.

10. An orbital energy-level diagram for sp2 hybridization.
09_164 2 p 2 p
Hybridization s p 2 E 2s p
Orbitals in an isolated
Carbon orbitals in ethylene
carbon atom
An orbital energy-level diagram for sp2 hybridization.

11. In sp2 hybridization one p orbital remains unchanged
and lies perpendicular to the plane of the hybrid.

12. The shared electron pair of in ethylene occupies
the region directly between the atoms to form a
sigma () bond.

13. sp2 hybrid orbital three effective electron pairs.
trigonal planar geometry.
120 o bond angle.

14. A carbon-carbon double bond consists of a  bond and a  bond
A carbon-carbon double bond consists of a  bond and a  bond. The  bond is formed from unhybridized p orbitals in the space above and below the  bond.

15. Pi and Sigma Bonds
Ϭ bonds consist of an electron pair shared in the area centered between the atoms.
Π bonds occupy the space above and below a line joining the atoms.

16. Pi and Sigma Bonds Ϭ bonds allow rotation.
Π bonds do not allow rotation.

17. The orbitals used to form the bonds in ethylene.

18. Two sp orbitals are formed when one s and one p orbital
are hybridized. They are oriented at 180o to each other.

19. The hybrid orbitals in the CO2 molecule.

20. sp hybrid orbital two effective electron pairs. linear geometry.
180 o bond angle.

21. The nitrogen molecule forms a triple bond -- one
 and two  bonds.

22. dsp3 hybrid orbitals five effective electron pairs.
trigonal bipyramidal geometry.
90 o and 120 o bond angles.
hybrid orbitals are not all equivalent as in the other types of hybridization.
Phosphorus pentachloride

23. d2sp3 hybrid orbitals six effective electron pairs.
octahedral geometry.
90 o bond angles.
Sulfur hexafluoride.

24. The relationship of the number of effective pairs,
their spatial arrangement, and the hybrid orbitals.

25. Counting Sigma and Pi Bonds
Type of Bond
# of Sigma Bonds
# of Pi Bonds
Single 1
Double
Triple 2

26. The Localized Electron Model
Draw the Lewis structure(s)
Determine the arrangement of electron pairs (VSEPR model).
Specify the necessary hybrid orbitals.

27. Deficiencies of the LEM Model
Does not adequately explain resonance.
Does not work for odd-electron molecules and ions.
Assumes that all electrons are localized about an atom.
Gives no direct information about bond energies.

28. Molecular Orbitals (MO)
Analagous to atomic orbitals for atoms, MOs are the quantum mechanical solutions to the organization of valence electrons in molecules. Electrons are considered to be delocalized over the entire molecule.

29. Types of MOs
bonding: lower in energy than the atomic orbitals from which it is composed.
antibonding: higher in energy than the atomic orbitals from which it is composed.

30. The molecular orbital energy diagram for the H2
molecule and the MO1 and MO2 orbitals formed.
MO1 = 1s and MO2 = 1s*.

31. Bond Order (BO)
Difference between the number of bonding electrons and number of antibonding electrons divided by two.
Larger bond order means greater bond strength!

32. The molecular orbital energy-level diagrams, bond orders,
bond energies, and bond lengths for diatomic molecules.

33. In order to participate in MOs, atomic orbitals must overlap in space
In order to participate in MOs, atomic orbitals must overlap in space. (Therefore, only valence orbitals of atoms contribute significantly to MOs.)

34. The relative size of the lithium 1s and 2s orbitals.
The 1s orbital can be considered to be localized
and do not participate in bonding.

35. The boron molecule will form one  and two 
bonds.

36. The two p orbitals that overlap head on make two 
molecular orbitals -- one bonding and one antibonding.
The two p orbitals that lie parallel overlap to produce two
 molecular orbitals, one bonding and one antibonding.

37. Paramagnetism unpaired electrons attracted to induced magnetic field
much stronger than diamagnetism
B2 & O2

38. The molecular orbital energy-level diagrams, bond orders,
bond energies, and bond lengths for diatomic molecules.

39. Diamagnetism paired electrons repelled from induced magnetic field
much weaker than paramagnetism
C2 , N2 , & F2 .

40. Apparatus used to measure the paramagnetism
of a sample. A paramagnetic sample will appear
heavier when the electromagnet is turned on.

41. A partial molecular orbital energy-level diagram
for the HF molecule. Bond order is 1 -- a single
bond.

42. The electron probability distribution in the bonding
molecular orbital of the HF molecule.

43. Pi and Sigma Bonds
 bonds in a molecule are described as being localized.
 bonds are considered to be delocalized over the entire molecule.

44. NO2 Molecule Draw the Lewis Structure, determine the
parent geometry, the actual geometry, and
the approximate bond angle.

45. NO+ ION Draw the Lewis Structure, draw the
molecular orbital energy-level diagram,
determine the bond order, and the type of
magnetism for the NO+ ion.

46. The  bonding system in the benzene molecule.

47. The  molecular orbital system in benzene. The
electrons in the  orbitals are delocalized over the
ring of carbon atoms.

48. Outcomes of MO Model
1. As bond order increases, bond energy increases and bond length decreases.
2. Bond order is not absolutely associated with a particular bond energy.
3. N2 has a triple bond, and a correspondingly high bond energy.
4. O2 is paramagnetic. This is predicted by the MO model, not by the LE model, which predicts diamagnetism.

49. Combining LE and MO Models
 bonds can be described as being localized.
 bonding must be treated as being delocalized.