1. Orbital Bonding Theory

2. Review:
Orbital: 3 dimensional space 1 or 2 electrons occupy around the nucleus of an atom
Basic Shapes:

3. Properties of Waves Frequency - the number of waves per second Wavelength - distance between each wave Amplitude - size of each wave above (+) or below (-) the axis Node – point where the amplitude = zero

4. + + + - - -
Node

5. Wave Interactions
Reinforcement and Interference of Waves Reinforcement: overlapping waves that are in phase (both + or both -)that increase the amplitude Interference: out of phase waves (one + and one -) that reduce the amplitude of both waves

8. Relationship to Bonding
For stable bonding to occur overlapping orbitals must be IN PHASE giving rise to reinforcement and a boding molecular orbital.
Out of phase orbitals result in interference and an antibonding molecular orbital.

9. Example: H + H = H2
Each atom contains one electron in the 1s orbital.
As the atoms approach each other the orbitals overlap in a reinforcing pattern and form a bonding molecular orbital that encompasses both nuclei and contains the two electrons (a shared pair)

11. The result of the reinforcing orbitals is an area of high electron density between the nuclei. This prevents the positively charged nuclei from repelling each other and the atoms stay bonded together.
This molecular orbital is cylindrically symmetrical around the axis (it can rotate without breaking) between the two hydrogen atoms and is called a sigma (σ) molecular orbital or a sigma (σ) bond.

12. Bonding of Out of Phase Orbitals
Out of phase orbitals, when they interact, produce a node of low electron density between the nuclei.
Since there is little negative charge due to the lack of electron density, the positive nuclei will repel each other resulting in a higher energy, less stable bond.
This is called an antibonding orbital or a sigma star (σ*) bond

14. Sigma bond formation in p sublevels
p sublevels can also form sigma bonds (and antisigma bonds) when their orbitals align and overlap in a horizontal axis

15. Carbon Orbital Hybridization
Carbon: 1s2 2s2 2p2
The 1st level electrons are not used for bonding because the energy level is full and stable.
The second energy level has four electrons, and thus, it seems, would be able to bond four times to fulfill the octet rule.
However, the two electrons in the 2s sublevel are fairly stable and would not normally bond.
Thus the orbitals can hybridize or blend the 2s orbital and the three 2p orbitals into sp orbitals

16. Carbon Orbital Hybrids
sp3: when carbon forms 4 single bonds
sp2: when carbon forms 1 double bond called a pi bond
sp: when a carbon forms 1 triple bond or 2 double bonds on the same carbon

17. sp3 bonding
Hybridization of the 2s orbital and the three 2p orbitals result in four equal orbitals called sp3 orbitals. This spreads the four valence electrons around the nucleus equally and allows for the carbon to bond four times forming the tetrahedron with the equivalent bond angles of 109.5o.

18. Area used for bonding

19. sp3 bonding
As the sp3 orbitals approach other orbitals (s, sp3), they will form sigma bonds when they overlap.
If with an s = sp3-s sigma bond
If with another sp3 = sp3-sp3 sigma bond
These are all single bonds and can rotate freely around the axis of the bond.

20. sp3 bonding in methane
The 1s orbitals of hydrogen overlap with the sp3 orbitals of the carbon to form 4 sp3-s sigma bonds.

21. sp3 bonding in ethane C2H6
Each carbon forms 3 sp3-s sigma bonds with hydrogen. The last bond of each carbon forms between the two carbons making one sp3-sp3 sigma bond.
sp3-s
sp3-sp3

22. sp2 bonding
In this state carbon hybridizes two of its p orbitals with the 2s orbital to create three sp2 hybrids.
These three orbitals spread out evenly and form a triangular plane with the bond angles = 120o.
This is known as a trigonal carbon or trigonal planar.

23. The remaining p orbital does not hybridize but exists perpendicular to the trigonal planar carbon sp2 orbitals.

24. If two sp2 carbons or another atom with a p orbital line up so the p orbitals form a double bond by overlapping side to side rather than end to end like in a sigma bond.
This bond is called a pi bond or a p bond.

25. Pi bond formation

26. Bonding in Ethene (C2H4)
Each of the two carbons form three sp2 orbitals and have 1 unhybridized p orbital.
Each carbon binds with two hydrogens forming a total of 4 sp2-s bonds.
The carbons bond to each other two ways:
1) the first bond is the sp2-sp2 sigma bond between the remaining sp2 orbitals
2) the second is the pi bond between the unhybridized p orbitals.

27. Bonding in Ethene
sp2-sp2 σ bond

28. Characteristics of pi bonds
Are not cylindrically symmetrical like sigma bonds.
-implication: these bonds cannot rotate or the bond will break.
- locks side groups into a particular geometry (cis and trans)

29. 2) pi bonds have more energy than sigma bonds and are less stable - implication: when a double bond breaks to become a single bond, it is the pi bond that breaks, not the sigma 3) pi bonds exist farther from the nucleus and are more exposed - implication: pi bonds are more “vulnerable” to attack by another atom or molecule making it the site of chemical reactivity = location of chemical reactions

30. Nucleophilic Addition Reaction

31. sp orbitals
Carbon hybridizes the 2s and one 2p orbital to form two sp orbitals.
The other 2 p orbitals remain unhybridized.
The sp orbitals will form sigma bonds and the 2 p orbitals will each form a pi bond
- Result: triple bond formed from one sigma and 2 pi bonds.

32. Bonding in Ethyne Ethyne = C2H2
Each carbon forms one bond with a hydrogen
Hydrogen s orbital and Carbon sp orbital = sp-s sigma bond.
Carbons bond to each other three times in two ways
1) sp-sp sigma bond
2) 2 pi bonds that are perpendicular to each other

34. Propadiene