1. Theories of Chemical Bonding
Exam #4 (Chapter 8) on 7-December
One 3” x 5” notecard
Chapter 8 OWL Deadline Tuesday Night
Chapter 9 OWL Final Exam
Final Exam (review session weekend before)
Monday, 8AM in IRC 3 (here)
~50 points just on Chapter 9
(about the same as the other chapters)
~100 points cumulative (Chapters 2-8)
One 81/2” x 11” note sheet allowed
Final Lab Experiment
This Week (let’s WIKI Page)
Laboratory Checkout Next Week!

2. What’s Gotten Us Here…. Electron Configurations Lewis Structures
Energy Level Diagrams
Valence Electrons
Lewis Structures
Identify Central Atom, Bonding and Lone Pairs
Formal Charge
Periodic Trends
Size, Electronegativity
Bond Properties
Bond Energy, Length, Order, Polarity

3. Electron Pair Geometry Molecular Geometry
Clues to the shape of the molecule or ion
Molecular Geometry
Definitive shape of the molecule or ion
Bond Angles…
However, other than identifying ionic and covalent bonds, we haven’t really described how atoms bond together!
Electrons in a covalent bond are shared, but how?

4. Let’s Look at Boron Trifluoride
B has 3 valence e-
Fluorine has 7 valence e-
We know that B follows an exception to the octet rule
Let’s look at the Lewis Structure

5. How do these 3 bonds form?
We have a single s and three p orbitals on the boron atom
We have a single s and three p orbitals on the fluorine atom
How do they overlap to share electrons?
Hybridization!

6. Valence Bond and Molecular Orbital Theories
Share some principles
Bonds are formed when electrons are shared between atoms
The sharing of electrons, and their attraction to the two nuclei of the atoms that are bonded lowers the total energy in the molecule or ion
Two types of bonds can form
sigma (σ) – bonds which lie along the axis between to atoms in a molecule or ion
Atoms can rotate around sigma bonds
pi (π) – bonds which lie in regions outside of the axis between two atoms (but parallel to the axis)
Atoms are fixed and can’t rotate
Multiple bonding

7. Valence Bond Theory Takes into account
Lewis structure (# of bonds)
VSEPR (overall shape of the molecule or ion)
The available orbitals that can be used for bonding
Quantum Mechanics
Spectroscopy results (like the colorimetry you used in lab)
Thermodynamic data
Helps decide what types of hybrids will form

8. Hybridization (very systematic)
To obtain the bonding description about any atom in a molecule:
1. Write the Lewis electron-dot formula.
2. Use VSEPR to determine the electron arrangement about the atom.
3. From the arrangement, deduce the hybrid orbitals.
4. Assign the valence electrons to the hybrid orbitals one at a time, pairing only when necessary.
5. Form bonds by overlapping singly occupied hybrid orbitals with singly occupied orbitals of another atom.

9. Some Tenets of Hybridization
The total number of orbitals you started with (all valence) MUST equal the number of orbitals that you end up with
The number of hybrid orbitals must equal the number of initial orbitals you started with (again all valence)
Two electrons per orbital
The number of sigma bonds equals the number of single bonds
Multiple bonds are sigma + pi bonds
s, p, d orbitals are what we are working with
Valence ONLY!!!

10. Let’s Start with Methane (CH4)
Write the valence electron configurations
Carbon has 4 valence electrons
4 hydrogen atoms each have 1 valence e-
Draw the Lewis structure
Evaluate electron and molecular structure
Shape helps us understand hybridization
Determine what orbitals will be combined (hybridized) to create the necessary bonds
Describe the hybridization

14. Lone Pairs ?
Lone pairs of electrons can be placed in hybrid orbitals too.
Ammonia
Write electron configurations (valence)
Lewis Structure
Predict electron and molecular geometry
Determine what orbitals might combine and hybridize to describe the bonding
Consider the best way to locate lone pairs (VSEPR)
Describe the hybridization
Let’s do this on the board!

16. OK, Let’s Get Back to BF3 Lewis Structure
Electron Pair and Molecular Structure (geometry)
Predict Hybridization
Remember-we are normally talking about hybridization about central atoms at this point.
Describe Hybridization

19. 3rd Period Elements (d orbitals)
Hybrids can be made from combinations of s, p and d orbitals.
Only available to elements in the 3rd period or higher (n = 2 only has s and p orbitals available)
Same rules apply
Total number of orbitals are constant
Number of hybrid orbitals equals the number of orbitals combined to make them
Some orbitals may be left unhybridized

20. 5 electron pairs (sp3d)

22. 6 electron pairs (sp3d2)

24. How is this systematic ?

25. Hybrid orbitals are named by using the atomic orbitals that combined:
• one s orbital + one p orbital gives two sp orbitals
• one s orbital + two p orbitals gives three sp2 orbitals
• one s orbital + three p orbitals gives four sp3 orbitals
• one s orbital + three p orbitals + one d orbital gives five sp3d orbitals
• one s orbital + three p orbitals + two d orbitals gives six sp3d2 orbitals

27. What about lone pairs (on the central atom)
We’ve seen NH3, where they go in one of the hybrid orbitals
Usually this is the case, because the hybrid orbital represents the greatest separation of electron pairs (bonding or lone) in 3D space.
They can go in other hybrid orbitals in other cases, for example…..

28. 4 total pairs (2 structural, 2 lone)
SeH2
Valence Electron Configuration
Lewis Dot Structure
Predict Electron Pair and Molecular Geometry
Different-two lone pairs.
Predict Hybridization
Draw with electron pairs in the appropriate orbitals and describe the hybridization
What happens to the tetrahedral bond angles?
Let’s do this on the board.

29. 5 total pairs (4 structural, one lone)

30. pi (π) bonds
sigma (σ) bonds lie along the internuclear axis between atoms
pi (π) bonds are parallel to the internuclear axis, but lie separated from it in space
Created from p orbitals not used in hybridization
Responsible for multiple bonds
We will concern ourselves with pi bonds in carbon containing (organic) molecules only

31. Double Bonds (one σ, one π)
Ethene (ethylene):
C2H4
3 x sp2 hybrid orbitals on each carbon
2 C-H sigma bonds on each carbon
1 C-C sigma bond
1 C-C pi bond
Double bond C-C (one sigma plus one pi)

33. Note that the geometry
of the sp2 hybrids (trigonal planar) controls geometry about each C atom.

35. Triple Bonds (one σ, two π)
Acetylene
C2H2
2 sp hybrid orbitals on each carbon
1 C-H sigma bonds on each carbon
1 C-C sigma bond
2 C-C pi bonds
Triple bond C-C (one sigma plus two pi)

37. Note that the geometry
of the sp hybrids (linear) controls the
geometry about each C atom.

38. What’s different about multiple bonds (sigma + pi) vs
What’s different about multiple bonds (sigma + pi) vs. single (sigma) bonds?
Location of the bonds
Along the internuclear axis versus parallel to it
Single bonds (sigma only) can rotate
Multiple bonds (sigma + pi) are rigid
Leads to structural isomers
cis vs. trans forms
More than one multiple (sigma + pi) in a molecule can lead to different resonance forms

39. Do we have cis vesus trans 1, 2 dichloroethane? Let’s draw on the board.

40. Benzene (C6H6) / Resonance
6 carbons, each with 3 sp2 hybrid orbitals
Each results in 3 sigma bonds
1 sigma bond to H, 2 sigma bonds to C
One pi bond on each carbon atom

43. sp2 hybridization still controls geometry on each carbon atom
pi bonds make the molecule “flat”
Two resonance forms describe possible location of pi bonds
How does this influence bond order, bond length, bond energy?
VERY stable-electrons are distributed throughout the molecule.

44. Molecular Orbital (MO) Theory
Considers orbitals to be more spread out (delocalized) than Valence Bond Theory
Spectroscopy results are more supportive of MO theory
Is better at describing why certain molecules do not form (why bonding does not happen in some cases)
Can be used in a complimentary fashion with Valence Bond Theory to describe what is happening in molecules.
We still have sigma (σ) and pi (π) bonding.
We add non-bonding and antibonding variables

45. How to conceptualize MO theory?
Instead of thinking of how hybrids form on a central atom….
Think about how orbitals on two separate atoms in a bond can come together and form bonds in a molecule
Combining orbitals (MO) instead of overlapping orbitals (VBT)

46. What is some of this terminology? What do the symbols mean?

47. Final Exam & Reminders
OWL Deadline (Chapter 9) is the start of the final exam
Review Sessions (IRC 3-here)
Sunday 10-11:30 AM
Sunday 12:30-2 PM
Final Exam (150 points)
45 points Chapter 9 (short answer)
128 points Chapters 2-8 (multiple choice)
Pen/Pencils, Calculator, Ruler, 81/2” x 11” note sheet
I will provide thermodynamic and other constants
Planck’s constant, speed of light, bond enthalpies, etc.

50. Some important principles of MO Theory
The number of molecular orbitals formed equals the number of atomic orbitals that have combined!
Electrons are spread out over the entire molecule
As opposed to being more localized in specific bonds as we see with VBT
The total energy of the molecular orbitals equals the sum of the energy of the atomic orbitals
It is just that some orbitals don’t contain electrons (antibonding)

51. What about He2?

52. H2 forms, He2 does not. Bond Order?
½ (# bonding electrons - # non-bonding electrons)
What is the bond order for H2 ?
What is the bond order for He2 ?
Bond orders less than zero suggest a molecule would not ordinarily exist.
A bond order between 0 and 1 suggests that the molecule would also not exist or would be unstable

53. Other homonuclear diatomic molecules (through F2)

54. Nitrogen (N2) Bond Order ? Diamagnetic or Paramagnetic ?

55. Oxygen (O2) Bond Order ? Diamagnetic or Paramagnetic ?

58. Heteronuclear Diatomic Molecules
The total energy of the orbitals created is still the same as the total of the atomic orbitals, but the energy of the original “incoming” atomic orbitals are different.
Same rules apply in terms of # of orbitals, magnetic properties, bond order, etc.

61. MO vs. VB Theories…..