1. Orbitals and Covalent Bond
Chapter 9
Orbitals and Covalent Bond

2. Molecular Orbitals
The overlap of atomic orbitals from separate atoms makes molecular orbitals
Each molecular orbital has room for two electrons
Two types of MO
Sigma (  ) between atoms
Pi (  ) above and below atoms

3. Sigma bonding orbitals
From s orbitals on separate atoms +
+ + + + +
Sigma bonding molecular orbital
s orbital
s orbital

4. Sigma bonding orbitals
From p orbitals on separate atoms  
p orbital
p orbital    
Sigma bonding molecular orbital

5. Pi bonding molecular orbital
Pi bonding orbitals
p orbitals on separate atoms        
Pi bonding molecular orbital

6. Sigma and pi bonds All single bonds are sigma bonds
A double bond is one sigma and one pi bond
A triple bond is one sigma and two pi bonds.

7. Atomic Orbitals Don’t Work
to explain molecular geometry.
In methane, CH4 , the shape is tetrahedral.
The valence electrons of carbon should be two in s, and two in p.
the p orbitals would have to be at right angles.
The atomic orbitals change when making a molecule

8. Hybridization
We blend the s and p orbitals of the valence electrons and end up with the tetrahedral geometry.
We combine one s orbital and 3 p orbitals.
sp3 hybridization has tetrahedral geometry.

11. In terms of energy 2p
Hybridization
sp3
Energy 2s

12. How we get to hybridization
We know the geometry from experiment.
We know the orbitals of the atom
hybridizing atomic orbitals can explain the geometry.
So if the geometry requires a tetrahedral shape, it is sp3 hybridized
This includes bent and trigonal pyramidal molecules because one of the sp3 lobes holds the lone pair.

13. sp2 hybridization C2H4 Double bond acts as one pair. trigonal planar
Have to end up with three blended orbitals.
Use one s and two p orbitals to make sp2 orbitals.
Leaves one p orbital perpendicular.

16. In terms of energy 2p 2p
Hybridization
sp2
Energy 2s

17. Where is the P orbital? Perpendicular
The overlap of orbitals makes a sigma bond (s bond)

18. Two types of Bonds Sigma bonds from overlap of orbitals.
Between the atoms.
Pi bond (p bond) above and below atoms
Between adjacent p orbitals.
The two bonds of a double bond.

19. H H C C H H

20. sp2 hybridization When three things come off atom. trigonal planar
120º
One p bond, s + lp =3

21. What about two When two things come off. One s and one p hybridize.
linear

22. sp hybridization End up with two lobes 180º apart.
p orbitals are at right angles
Makes room for two p bonds and two sigma bonds.
A triple bond or two double bonds.

23. In terms of energy 2p 2p sp
Hybridization
Energy 2s

24. CO2
C can make two s and two p
O can make one s and one p O C O

25. N2

26. N2

27. Breaking the octet
PCl5
The model predicts that we must use the d orbitals.
dsp3 hybridization
There is some controversy about how involved the d orbitals are.

28. dsp3 Trigonal bipyrimidal can only s bond. can’t p bond.
basic shape for five things.

29. PCl5 Can’t tell the hybridization of Cl
Assume sp3 to minimize repulsion of electron pairs.

30. d2sp3
gets us to six things around
Octahedral
Only σ bond

31. Molecular Orbital Model
Localized Model we have learned explains much about bonding.
It doesn’t deal well with the ideal of resonance, unpaired electrons, and bond energy.
The MO model is a parallel of the atomic orbital, using quantum mechanics.
Each MO can hold two electrons with opposite spins
Square of wave function tells probability

32. What do you get? Solve the equations for H2 HA HB get two orbitals
MO2 = 1sA - 1sB
MO1 = 1sA + 1sB

33. The Molecular Orbital Model
The molecular orbitals are centered on a line through the nuclei
MO1 the greatest probability is between the nuclei
MO2 it is on either side of the nuclei
this shape is called a sigma molecular orbital

34. The Molecular Orbital Model
In the molecule only the molecular orbitals exist, the atomic orbitals are gone
MO1 is lower in energy than the 1s orbitals they came from.
This favors molecule formation
Called an bonding orbital
MO2 is higher in energy
This goes against bonding
antibonding orbital

35. The Molecular Orbital Model
Energy 1s 1s
MO1

36. The Molecular Orbital Model
We use labels to indicate shapes, and whether the MO’s are bonding or antibonding.
MO1 = s1s
MO2 = s1s* (* indicates antibonding)
Can write them the same way as atomic orbitals
H2 = s1s2

37. The Molecular Orbital Model
Each MO can hold two electrons, but they must have opposite spins
Orbitals are conserved.
The number of molecular orbitals must equal the number atomic orbitals that are used to make them.

38. H2-
s1s*
Energy 1s 1s
s1s

39. Bond Order
The difference between the number of bonding electrons and the number of antibonding electrons divided by two

40. Only outer orbitals bond
The 1s orbital is much smaller than the 2s orbital
When only the 2s orbitals are involved in bonding
Don’t use the s1s or s1s* for Li2
Li2 = (s2s)2
In order to participate in bonds the orbitals must overlap in space.

41. Bonding in Homonuclear Diatomic Molecules
Need to use Homonuclear so that we know the relative energies.
Li2-
(s2s)2 (s2s*)1
Be2
(s2s)2 (s2s*)2
What about the p orbitals? How do they form orbitals?
Remember that orbitals must be conserved.

42. B2

43. B2
s2p*
s2p
p2p*
p2p

44. Expected Energy Diagram
s2p*
p2p*
p2p* 2p 2p
p2p
p2p
s2p
Energy
s2s* 2s 2s
s2s

45. B2 2p 2p
Energy 2s 2s

46. B2 (s2s)2(s2s*)2 (s2p)2 Bond order = (4-2) / 2 Should be stable.
This assumes there is no interaction between the s and p orbitals.
Hard to believe since they overlap
proof comes from magnetism.

47. Magnetism Magnetism has to do with electrons.
Remember that spin is how an electron reacts to a magnetic field
Paramagnetism attracted by a magnet.
associated with unpaired electrons.
Diamagnetism repelled by a magnet.
associated with paired electrons.
B2 is paramagnetic.

48. Magnetism
The energies of of the p2p and the s2p are reversed by p and s interacting
The s2s and the s2s* are no longer equally spaced.
Here’s what it looks like.

49. Correct energy diagram
s2p*
p2p*
p2p* 2p
s2p 2p
p2p
p2p
s2s* 2s 2s
s2s

50. B2
s2p*
p2p* 2p 2p
s2p
p2p
s2s* 2s 2s
s2s

51. Patterns As bond order increases, bond energy increases.
As bond order increases, bond length decreases.
Supports basis of MO model.
There is not a direct correlation of bond order to bond energy.
O2 is known to be paramagnetic.
Movie.

52. Magnetism Ferromagnetic strongly attracted
Paramagnetic weakly attracted
Liquid Oxygen
Diamagnetic weakly repelled
Graphite
Water Frog

53. Examples C2 N2 O2 F2 P2

54. Heteronuclear Diatomic Species
Simple type has them in the same energy level, so can use the orbitals we already know.
Slight energy differences. NO

55. NO 2p 2p 2s 2s

56. You try
NO+
CN-
What if they come from completely different orbitals and energy? HF
Simplify first by assuming that F only uses one if its 2p orbitals.
F holds onto its electrons, so they have low energy

57. s* 1s 2p s

58. Consequences Paramagnetic
Since 2p is lower in energy, favored by electrons.
Electrons spend time closer to fluorine.
Compatible with polarity and electronegativity.

59. Names sp orbitals are called the Localized electron model
s and p Molecular orbital model
Localized is good for geometry, doesn’t deal well with resonance.
seeing s bonds as localized works well
It is the p bonds in the resonance structures that can move.

60. p delocalized bonding
C6H6 H H

61. C2H6

62. NO3-