1. Covalent Bonding Pt 3: Hybridization

2. Introduction
We now know that atoms can bond covalently through the sharing of electrons
VSEPR theory helps us predict molecular shapes. But, it does not explain what bonds are, how they form, or why they exist.
Here, chemical bonding will be explained in terms of orbitals

3. Covalent Bonding Is Due to Orbital Overlap
In a covalent bond, electron density is concentrated between the nuclei.
Thus, we can imagine the valence orbitals of the atoms overlapping
The region of orbital overlap represents the covalent bond

4. Overlapping Valence Orbitals
Recall s and p orbitals S
S orbitals are spherical. L = 0, mL = 0
Max of 2 electrons
P orbitals consist of two lobes of electron density.
L= 1, mL = -1, 0, 1
(3 suborbitals)
Max of 6 electrons px py pz

5. Forming Sigma (σ) Bonds
Energy
1s1
1s1 σ
stabilization (energy drop)
Two overlapping atomic orbitals form a molecular bonding orbital.
A sigma (σ) bonding orbital forms when s-orbitals overlap.

6. Introduction to Hybridization
Imagine the molecule CH4. We know that carbon has 4 valence electrons (2s22p2).
However, when we fill our orbitals in order as according to Hund’s rule, we notice there are only enough unpaired electrons to make two bonds. Stay mindful of the fact that a covalent bond involves the sharing of unpaired electrons C
ENERGY X
2p2
4 H
2s2
1s1
1s1
1s1
1s1

7. sp3 Hybridization So how does CH4 form? How can carbon make 4 bonds?
To make four bonds, carbon hybridizes four of its atomic orbitals. This creates four equivalent sp3 hybrid orbitals, each containing one unpaired electron.
The name “sp3” originates from the fact that the hybrid orbitals form as a result of the mixture of 1 s-orbital and 3 p-orbitals. Thus, each sp3 orbital is 25% s character and 75% p character

8. Illustration of sp3 Hybrid Orbitals and Orbital Overlap
Orbital Mixing
The four hybrid orbitals arrange themselves tetrahedrally around the nucleus.

9. Formation of Sigma Bonding Orbitals
sp3 hybrid orbitals
ENERGY C
1s1
1s1
1s1
1s1
atomic s-orbitals 4H
σ bonding orbitals

10. Hybridization of Lone Electron Pairs
Ex. What is the hybridization of Oxygen in H2O?
The valence electron configuration of O is [He]2s2 2p4 O
ENERGY
2p4
2 H
2s2
1s1
1s1
As you see, there are two unpaired O electrons. Does this mean that these two p-suborbitals can overlap with the two Hydrogen 1s orbitals without hybridizing??

11. •• •• X Hybridization of Lone Electron Pairs O 2 H O H H ENERGY 2p4
2p electrons
2s electrons O H H
ENERGY
2p4 X
2 H
BAD!!
2s2
1s1
1s1
No!! The reason is that we now have two sets of lone pairs of electrons that are substantially different in energy (2s and 2p). The orbitals will hybridize to form degenerate (equal energy) lone pairs.

12. Hybridization of Lone Electron Pairs
The orbitals will hybridize to form degenerate (equal energy) lone pairs.

13. A Simple Trick To Remember
Total Electron Domains Around Atom (Bond + LP)
Hybridization 2 sp 3
sp2 4
sp3

14. Double and Triple Bonding
How can orbital overlap be used to explain double and triple bonds? What kind of interactions are these?
Lets look at formaldehyde, H2CO C O H ••
sp2
The hybridization of carbon and oxygen is sp2 because each is surrounded by three electron domains.
sp2

15. Forming Double Bonds
We can see that the carbon atom needs to make three sigma bonds, so this will require three unpaired electrons to be spread across the three sp2 hybrid orbitals. This will leave one unhybridized p orbital with an electron in it.

16. Forming Double Bonds
The oxygen is also sp2. It needs to form two lone pair, and also needs to form one sigma bond to carbon. This will require 5 electrons to be spread across the three sp2 hybrid orbitals, which will leave one unhybridized p orbital with an electron in it.

17. Forming Double Bonds
The H s-orbitals overlap with C sp2 orbitals to create C-H sigma bonds.
The remaining C sp2 orbital overlaps with the half filled O sp2 orbital to create the C-O sigma bond.
The unhybridized p orbitals of C and O interact to produce a pi bond, which is slightly weaker than a sigma.
DOUBLE BOND = 𝟏𝝈+𝟏𝝅

18. Triple Bonds :1 σ-bond and 2 π-bonds. Ex. C2H2sp sp H C C H
The carbons have sp hybridization. Each carbon requires two electrons to make two sigma bonds. This leaves two unhybridized p orbitals, each containing one electron.

19. Examples CH3CH2CHCHCH3 CH3CCCHCH2
How many σ and π bonds are in each of the following molecules? Give the hybridization of each carbon.
CH3CH2CHCHCH3
CH3CCCHCH2

20. sp3d and sp3d2 hybridization
Atoms like S, Se, I, Xe … etc. can exceed an octet because of sp3d and sp3d2 hybridization (combination of ns, np, and nd orbitals where n>3).
This results in either trigonal bipyramidal or octahedral skeletal geometry
sp3d
sp3d2

21. Exceeding an Octet. Example: SF6
S needs to make six sigma bonds. Therefore, its 6 valence electrons must be spread across six hybrid orbitals. This means that the 3s orbital, all three of the 3p orbitals, and two of the 3d orbitals must mix to yield six sp3d2 hybrid orbitals. Note: The fluorines are sp3 hybridized.

22. Look Familiar ???

23. Examples: What is the hybridization of the central atom? CO2 H2CO
CH3CCH
IF5
PCl5
SeOF4