1. Chapter 11 Theories of Covalent Bonding
Hybrid and MO
4/23/2017
Chapter 11
Theories of Covalent Bonding
If you are doing this lecture “online” then print the lecture notes available as a word document, go through this ppt lecture, and do all the example and practice assignments for discussion time.

2. VALENCE BOND THEORY:
DEVELOPED BY LINUS PAULING, who received the Nobel Prize in 1954 for his work
A view of chemical bonding in which bonds arise from the overlap of atomic orbitals on two atoms to give a bonding orbital of electrons localized between the bonded atoms
RULE: Realize that Valence Bond Theory and all the others don't explain everything

3. The Central Themes of VB Theory
Basic Principle
A covalent bond forms when the orbitals of two atoms overlap and the overlap region, which is between the nuclei, is occupied by a pair of electrons.
(The two wave functions are in phase so the amplitude increases
between the nuclei.)

4. The Central Themes of VB Theory
A set of overlapping orbitals has a maximum of two electrons that must have opposite spins.
The greater the orbital overlap, the stronger (more stable) the bond.
The valence atomic orbitals in a molecule are different from those in isolated atoms.
There is a hybridization of atomic orbitals to form molecular
orbitals.

5. Orbital overlap and spin pairing in three diatomic molecules.
Figure 11.1
Hydrogen, H2
Hydrogen fluoride, HF
Regular atomic orbital overlap can explain these bonds.
Fluorine, F2

6. VALENCE BOND THEORY: Ha to Hb: 1sa to 1sb overlap radius = 74 pm
As overlap increases, strength of bond increases - both electrons are mutually attracted to both atomic nuclei.
At optimum distance between nuclei with maximum overlap, a sigma s bond (strong primary bond) forms. Max electron density is along the axis of the bond
Ha to Fb: 1sa to 2pb direct overlap or s bond

7. VALENCE BOND THEORY:
F to F: the picture looks like a 2p orbital on one F is overlapping with a 2p orbital on the other F atom, but actually each F is sp3 hybridized & electrons are localized between two atomic nuclei

8. VALENCE BOND THEORY:
We cannot use this direct overlap picture for CH4’s bonding. The 2s and the three 2p orbitals on each C do not fit into the CH4 molecule's 109o bond angles, since the 2p orbitals are at 90° to each other
Valence Bond Theory states that HYBRID orbitals of the outermost orbitals on an atom are formed from the atoms’ atomic orbitals

9. Hybrid Orbitals Key Points
The number of hybrid orbitals obtained equals the number of atomic orbitals mixed.
The type of hybrid orbitals obtained varies with the types of atomic orbitals mixed.
Types of Hybrid Orbitals sp
sp2
sp3
sp3d
sp3d2

10. The sp hybrid orbitals in gaseous BeCl2.
Figure 11.2
The sp hybrid orbitals in gaseous BeCl2.
atomic orbitals on Be
hybrid orbitals
You have to know how to draw this energy hybrid formation.
orbital box diagrams

11. The sp hybrid orbitals in gaseous BeCl2(continued).
Figure 11.2
The sp hybrid orbitals in gaseous BeCl2(continued).
orbital box diagrams with orbital contours

12. The sp2 hybrid orbitals in BF3.
Figure 11.3
The sp2 hybrid orbitals in BF3.
You have to know how to draw this energy hybrid formation.
Note the three sigma bonds formed between B and each F.

13. The sp3 hybrid orbitals in CH4.
Figure 11.4
The sp3 hybrid orbitals in CH4.
You have to know how to draw this energy hybrid formation.

14. The sp3 hybrid orbitals in NH3.
Figure 11.5
The sp3 hybrid orbitals in NH3.
You have to know how to draw this energy hybrid formation.

15. The sp3 hybrid orbitals in H2O.
Figure 11.5 continued
The sp3 hybrid orbitals in H2O.
You have to know how to draw this energy hybrid formation.

16. VALENCE BOND THEORY
Expanded Valence Shells have hybrid orbitals using s, p & d atomic orbitals. Example: PCl5 P: [Ne]3s23p3
dsp3 hybridization results in 5 s bonds and trigonal bipyramidal geometry
(You can write these as dsp3 or sp3d)

17. The sp3d hybrid orbitals in PCl5.
Figure 11.6
The sp3d hybrid orbitals in PCl5.
You have to know how to draw this energy hybrid formation.

18. The sp3d2 hybrid orbitals in SF6.
Figure 11.7
The sp3d2 hybrid orbitals in SF6.
You have to know how to draw this energy hybrid formation.

20. Molecular shape and e- group arrangement
Figure 11.8
The conceptual steps from molecular formula to the hybrid orbitals used in bonding.
Figure 10.1
Step 1
Figure 10.12
Step 2
Step 3
Table 11.1
Molecular shape and e- group arrangement
Molecular formula
Lewis structure
Hybrid orbitals

21. SAMPLE PROBLEM 11.1
Postulating Hybrid Orbitals in a Molecule
PROBLEM:
Use partial orbital diagrams to describe mixing of the atomic orbitals of the central atom leads to hybrid orbitals in each of the following:
(a) Methanol, CH3OH
(b) Sulfur tetrafluoride, SF4
SOLUTION:
The groups around C are arranged as a tetrahedron.
(a) CH3OH
O also has a tetrahedral arrangement with 2 nonbonding e- pairs.

22. SAMPLE PROBLEM 11.1
Postulating Hybrid Orbitals in a Molecule
continued
single C atom
single O atom
hybridized C atom
hybridized O atom
(b) SF4 has a seesaw shape with 4 bonding and 1 nonbonding e- pairs.
S atom
hybridized S atom

23. VALENCE BOND THEORY
There can be more than one central atom, and each has its own hybridization and geometry
C2H6 and H2O2 and CH3COOH
C2H6: both C's are sp3 hybridized and can rotate around axis of bond.
H2O2: both O's are sp3 , etc.

24. relatively even distribution of electron density over all s bonds
Figure 11.9
The s bonds in ethane(C2H6).
both C are sp3 hybridized
s-sp3 overlaps to s bonds
sp3-sp3 overlap to form a s bond
relatively even distribution of electron density over all s bonds

25. VALENCE BOND THEORY: Multiple Bonds
H2CO: the Lewis structures shows a double bond between C and O, but we know it does not have twice the bond dissociation energy of a single C-O bond
Pauling proposed that there was only one sigma s bond between any two atoms, and the other multiples were weaker pi p bonds
If there are only 3 s bonds around this carbon, it can't be sp3 hybridized - instead we have sp2 hybrid orbitals
sp2 hybridization results in only 3 s bonds, and trig planar geometry, with 120° angles
p bond is a sideways or parallel overlap of the p atomic orbitals rather than the direct overlap of s bonds

26. The s and p bonds in ethylene (C2H4).
Figure 11.10
The s and p bonds in ethylene (C2H4).
overlap in one position - s
p overlap - 
electron density
Proper name is ethene.

27. VALENCE BOND THEORY
Look at acetylene: its geometry is linear. C is forming a triple bond to another C and a single bond to H, so that's only two s bonds
Therefore sp hybridization results in only 2 s bonds, and linear geometry
There are 2 p bonds from the parallel overlap of the 2p orbitals remaining on both C's

28. The s and p bonds in acetylene (C2H2).
Figure 11.11
The s and p bonds in acetylene (C2H2).
overlap in one position - s
p overlap - 

29. SAMPLE PROBLEM 11.2
Describing the Bond in Molecules
PROBLEM:
Describe the types of bonds and orbitals in acetone, (CH3)2CO.
PLAN:
Use the Lewis structures to ascertain the arrangement of groups and shape at each central atom. Postulate the hybrid orbitals taking note of the multiple bonds and their orbital overlaps.
SOLUTION:
sp3 hybridized
sp2 hybridized
bond
bonds

30. H2CO hybrid orbitals and sigma and pi bond formation
Remember the C=C double bond has sigman and pi bonds.
The C has sigma bonds from its hybrid orbitals to the two H’s and the O. The leftover p orbitals will form the pi bond.

31. TRANS CIS Restricted rotation of p-bonded molecules in C2H2Cl2.
Figure from 4th ed.
Restricted rotation of p-bonded molecules in C2H2Cl2.
TRANS
CIS
This cis/trans arrangement will be important in chem 2, organic chem and biology!

32. VALENCE BOND THEORY: RESONANCE
Resonance Structures and p Bonding:
p resonance structures involve an electron pair used alternately as a p bond or a LP
Ozone: O3 O==O--O or O--O==O
All are sp2, trig planar, each has 3 sp2 orbitals and a p orbital remaining.

33. VALENCE BOND THEORY
Benzene: C6H6 has carbons with sp2 hybrids and 120o angles, each C has 2 s bonds to other C's, 1 s bond to H, and 1 p bond electron available
Get "ring" of delocalized p e-s
SUMMARY: draw the Lewis structure; determine arrangement of electron pairs using VSEPR, specify the hybrid orbitals to accommodate the e- pairs

34. Benzene sigma bond formation between C’s and C-Hs
Hybrid and MO
4/23/2017
Benzene sigma bond formation between C’s and C-Hs
The leftover p orbitals will form alternating pi bonds as shown in sketch.

35. MOLECULAR ORBITAL THEORY:
- explains why H2 forms easily and He2 does not
- is an alternate way of viewing e- orbitals in molecules where pure s and pure p orbitals combine to produce orbitals that are delocalized over the molecule
- they can have different energies and are assigned electrons just like we do in an atom - Pauli exclusion principle and Hund's rule included
Pauling's Valence Bond Theory does not explain everything
MO Theory doesn't either, but it does correctly predict the electronic structure of certain molecules that do not follow Lewis's approach, including the paramagnetism of certain molecules, like O2

36. The Central Themes of MO Theory
A molecule is viewed on a quantum mechanical level as a collection of nuclei surrounded by delocalized molecular orbitals.
Atomic wave functions are summed to obtain molecular wave functions. The number of molecular orbitals produced is always = # of atomic orbitals brought by the combining atoms (only orbitals on different atoms are combined).
If wave functions reinforce each other, a bonding MO is formed (region of high electron density exists between the nuclei).
If wave functions cancel each other, an antibonding MO is formed (a node of zero electron density occurs between the nuclei).
The electrons of the molecule are placed in bonding or antibonding orbitals of successively higher energy (just like Hund's rule).
Atomic orbitals combine most effectively with orbitals of the same type and similar energy (s w/s, n=2 w/ n=2)

37. Amplitudes of wave functions subtracted.
Figure 11.13
An analogy between light waves and atomic wave functions.
Amplitudes of wave functions subtracted.
Amplitudes of wave functions added

38. Figure 11.14
Contours and energies of the bonding and antibonding molecular orbitals (MOs) in H2.
The bonding MO is lower in energy and the antibonding MO is higher in energy than the AOs that combined to form them.

39. MOLECULAR ORBITAL THEORY
BOND ORDER: the number of bonding e- pairs shared by 2 atoms in a molecule
Fractional bond orders are possible in MO Theory!
Silberberg method:
B.O. = ½(# of e- in bonding orbitals - # of e- in antibonding orbitals)

40. The MO diagram for H2. Energy AO of H AO of H
Figure 11.15
The MO diagram for H2.
Filling molecular orbitals with electrons follows the same concept as filling atomic orbitals.
s*1s
Energy
AO of H 1s
AO of H 1s
H2 bond order = 1/2(2-0) = 1
s1s
MO of H2

41. MO diagram for He2+ and He2.
Figure 11.16
MO diagram for He2+ and He2.
Energy
s*1s
s1s
s*1s
AO of He 1s
AO of He+ 1s
AO of He 1s
AO of He 1s
Energy
s1s
MO of He+
MO of He2
He2 bond order = 0
He2+ bond order = 1/2

42. SAMPLE PROBLEM 11.3
Predicting Stability of Species Using MO Diagrams
PROBLEM:
Use MO diagrams to predict whether H2+ and H2- exist. Determine their bond orders and electron configurations.
SOLUTION:
bond order = 1/2(1-0) = 1/2
bond order = 1/2(2-1) = 1/2 s s s s
H2+ does exist
H2- does exist 1s
AO of H
AO of H- 1s
AO of H 1s
AO of H
MO of H2-
MO of H2+

43. Bonding in s-block homonuclear diatomic molecules.
Figure 11.17
Bonding in s-block homonuclear diatomic molecules.
s*2s
s2s
s*2s
s2s
Energy 2s 2s 2s
Be2
Li2
Be2 bond order = 0
Li2 bond order = 1

44. Contours and energies of s and p MOs through combinations of 2p atomic orbitals.
Figure 11.18
Or the pz orbitals

45. Relative MO energy levels for Period 2 homonuclear diatomic molecules.
Figure 11.19
Relative MO energy levels for Period 2 homonuclear
diatomic molecules.
with 2s-2p mixing
without 2s-2p mixing
Memorize this!
MO energy levels for O2, F2, and Ne2
MO energy levels for B2, C2, and N2

46. MO occupancy and molecular properties for B2 through Ne2
Figure 11.20
MO occupancy and molecular properties for B2 through Ne2

47. The paramagnetic properties of O2
Figure 11.21
The paramagnetic properties of O2

48. SAMPLE PROBLEM 11.4
Using MO Theory to Explain Bond Properties
PROBLEM:
As the following data show, removing an electron from N2 forms an ion with a weaker, longer bond than in the parent molecules, whereas the ion formed from O2 has a stronger, shorter bond:
Bond energy (kJ/mol)
Bond length (pm) N2
N2+ O2
O2+
945
110
498
841
623
112
121
Explain these facts with diagrams that show the sequence and occupancy of MOs.
SOLUTION:
N2 has 10 valence electrons, so N2+ has 9.
O2 has 12 valence electrons, so O2+ has 11.

49. SAMPLE PROBLEM 11.4
Using MO Theory to Explain Bond Properties
continued N2
N2+ O2
O2+
2p
s2s
s2s
2p
2p
2p
2p
antibonding e- lost
bonding e- lost
2p
2p
2p
s2s
s2s
bond orders
1/2(8-2)=3
1/2(7-2)=2.5
1/2(8-4)=2
1/2(8-3)=2.5

50. MO Theory Practice
1. Draw the bonding and antibonding molecular orbitals for H2.
2. Do Valence Bond Theory (hybridization) and MO Theory for both O2 and O22-. Which theory works better to explain the molecule and ion?
3. For N2, N2+ and N2- compare
a. Magnetic character
b. Net number of p bonds
c. Bond Order
d. Bond length
e. Bond strength

51. Answers 1. See picture in text.
2. VB Theory shows O2 as sp2 hybridized with one s bond and one p bond. There are two lone pairs on each O. O22- has one s bond, and each O has three lone pairs. MO Theory shows a bond order of 2 for O2 and that it is paramagnetic. MO Theory shows a bond order of 1 for O22- and diamagnetic. But MO Theory fits the real data that O2 is paramagnetic.

52. Answers con’t N2 N2+ N2- a. Diamag Paramag b. 2 1.5 c. 3 2.5 d. Short
Longer e.
Strong
Weaker