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 sp 3 hybrid orbitals.

5. One 2s and three 2p orbitals hybridize to form a new set of sp 3 hybrid orbitals.

6. sp 3 hybrid orbital 4 effective electron pairs. tetrahedral geometry. 109.5 o bond angle.

7. The tetrahedral set of four sp 3 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 sp 3 hybridized.

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

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

11. In sp 2 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. sp 2 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. The  bond is formed from unhybridized p orbitals in the space above and below the  bond.

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

16. Pi and Sigma Bonds  bonds allow rotation.  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 180 o to each other.

19. The hybrid orbitals in the CO 2 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. dsp 3 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. d 2 sp 3 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. The Localized Electron Model -Draw the Lewis structure(s) -Determine the arrangement of electron pairs (VSEPR model). -Specify the necessary hybrid orbitals.

26. 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.

27. 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.

28. 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.

29. The molecular orbital energy diagram for the H 2 molecule and the MO 1 and MO 2 orbitals formed. MO 1 =  1s and MO 2 =  1s *.

30. 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!

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

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

33. 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.

34. The boron molecule will form one  and two  bonds.

35. 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.

36. Paramagnetism -unpaired electrons -attracted to induced magnetic field -much stronger than diamagnetism -B 2 & O 2

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

38. Diamagnetism -paired electrons -repelled from induced magnetic field -much weaker than paramagnetism -C 2, N 2, & F 2.

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

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

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

42. Pi and Sigma Bonds  bonds in a molecule are described as being localized.  bonds are considered to be delocalized over the entire molecule.

43. NO 2 Molecule Draw the Lewis Structure, determine the parent geometry, the actual geometry, and the approximate bond angle.

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

45. The  bonding system in the benzene molecule.

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

47. 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.N 2 has a triple bond, and a correspondingly high bond energy. 4.O 2 is paramagnetic. This is predicted by the MO model, not by the LE model, which predicts diamagnetism.

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