10. Valence Bond Theory

11. Orbital shapes, Individual (“isolated”) Atoms

12. 2s These new orbitals are called hybrid orbitals The process is called hybridization What this means is that both the s and one p orbital are involved in bonding to the connecting atoms Formation of sp hybrid orbitals The combination of an s orbital and a p orbital produces 2 new orbitals called sp orbitals.

13. Formation of sp 2 hybrid orbitals

14. Formation of sp 3 hybrid orbitals

15. To rationalize how the shapes of atomic orbitals are transformed into the orbitals occupied in covalently bonded species, we need the help of two bonding theories: Valence Bond (VB) Theory, the theory we will explore, describes the placement of electrons into bonding orbitals located around the individual atoms from which they originated. Molecular Orbital (MO) Theory places all electrons from atoms involved into molecular orbitals spread out over the entire species. This theory works well for excited species, and molecules like O 2. You will meet this theory in advanced classes!

16. COVALENT BOND FORMATION (VB THEORY) In order for a covalent bond to form between two atoms, overlap must occur between the orbitals containing the valence electrons. The best overlap occurs when two orbitals are allowed to meet “head on” in a straight line. When this occurs, the atomic orbitals merge to form a single bonding orbital and a “single bond” is formed, called a sigma (  ) bond.

18. MAXIMIZING BOND FORMATION In order for “best overlap” to occur, valence electrons need to be re-oriented and electron clouds reshaped to allow optimum contact. To form as many bonds as possible from the available valence electrons, sometimes separation of electron pairs must also occur. We describe the transformation process as “orbital hybridization” and we focus on the central atom in the species...

19. “sp” Hybridization: all 2 Region Species

20. Hybridization of Be in BeCl 2 Atomic Be: 1s 2 2s 2 Valence e’s Hybrid sp orbitals: 1 part s, 1 part p

21. FORMATION OF BeCl 2 : Each Chlorine atom, 1s 2 2s 2 2p 6 3s 2 3p 5, has one unshared electron in a p orbital. The half filled p orbital overlaps head-on with a half full hybrid sp orbital of the beryllium to form a sigma bond.

23. “sp 2 ” Hybridization: All 3 Region Species

24. Hybridization of B in BF 3 Atomic B : 1s 2 2s 2 2p 1 Valence e’s Hybrid sp 2 orbitals: 1 part s, 2 parts p

25. FORMATION OF BF 3 : Each fluorine atom, 1s 2 2s 2 2p 5, has one unshared electron in a p orbital. The half filled p orbital overlaps head-on with a half full hybrid sp 2 orbital of the boron to form a sigma bond.

27. “sp 3 ” Hybridization: All 4 Region Species

28. Hybridization of C in CH 4 Atomic C : 1s 2 2s 2 2p 2 Valence e’s Hybrid sp 3 orbitals: 1 part s, 3 parts p

29. FORMATION OF CH 4 : Each hydrogen atom, 1s 1, has one unshared electron in an s orbital. The half filled s orbital overlaps head-on with a half full hybrid sp 3 orbital of the carbon to form a sigma bond.

31. Unshared Pairs, Double or Triple Bonds Unshared pairs occupy a hybridized orbital the same as bonded pairs: See the example of NH 3 that follows. Double and triple bonds are formed from electrons left behind and unused in p orbitals. Since all multiple bonds are formed on top of sigma bonds, the hybridization of the single (  ) bonds determine the hybridization and shape of the molecule...

33. Hybridization of N in NH 3 Atomic N: 1s 2 2s 2 2p 3 Valence e’s

34. FORMATION OF NH 3 : Each hydrogen atom, 1s 1, has one unshared electron in an s orbital. The half filled s orbital overlaps head-on with a half full hybrid sp 3 orbital of the nitrogen to form a sigma bond.

36. Group Work 13.1 Describe Hybridization of C and shape of following species: CO, CO 2, HCN, CH 2 O, CO 3 2-, CBr 4

37. “sp 3 d” Hybridization: All 5 Region Species

38. Hybridization of P in PF 5 P: 1s 2 2s 2 2p 6 3s 2 3p 3

39. FORMATION OF PF 5 : Each fluorine atom, 1s 2 2s 2 2p 5, has one unshared electron in a p orbital. The half filled p orbital overlaps head-on with a half full hybrid sp 3 d orbital of the phosphorus to form a sigma bond.

41. “sp 3 d 2 ” Hybridization: All 6 Region Species

42. Hybridization of S in SF 6 S: 1s 2 2s 2 2p 6 3s 2 3p 4

43. FORMATION OF SF 6 : Each fluorine atom, 1s 2 2s 2 2p 5, has one unshared electron in a p orbital. The half filled p orbital overlaps head-on with a half full hybrid sp 3 d 2 orbital of the phosphorus to form a sigma bond.

45. Group Work 13.2 Describe hybridization of S and shape of species in SF 2, SO 2, SO 3 2-, SF 3 +, SF 4, SF 5 -

46. Summary: Regions, Shapes and Hybridization

47. Multiple Bonds Everything we have talked about so far has only dealt with what we call sigma bonds Sigma bondA bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms. Sigma bond (  )  A bond where the line of electron density is concentrated symmetrically along the line connecting the two atoms.

48. Pi bondA bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis). Pi bond (  )  A bond where the overlapping regions exist above and below the internuclear axis (with a nodal plane along the internuclear axis).

49. Example: H 2 C=CH 2

51. Example: HC  CH

52. Delocalized  bonds When a molecule has two or more resonance structures, the pi electrons can be delocalized over all the atoms that have pi bond overlap.

53. In general delocalized  bonding is present in all molecules where we can draw resonance structures with the multiple bonds located in different places. Benzene is an excellent example. For benzene the  orbitals all overlap leading to a very delocalized electron system Example: C 6 H 6 benzene

54. The dipole moment (  )  = Qr where Q is the magnitude of the charges and r is the distance

55. The sum of these vectors will give us the dipole for the molecule For a polyatomic molecule we treat the dipoles as 3D vectors