1. Qantum Mechanics and Bonding Hybridization
4.6

2. Bonding
Atoms form covalent bonds when the orbitals overlap. These orbitals are now called molecular orbitals.

3. Let’s consider carbon…
How many valence electrons?
In which orbitals?
So, both the 2s and 2p orbitals are used to form bonds
How many bonds does carbon form?
All four C-H bonds are the same
i.e. there are not two types of bonds from the two different orbitals
How do we explain this?
Hybridization

4. Hybridization
The s and p orbitals of the C atom combine with each other to form hybrid orbitals before they combine with orbitals of another atom to form a covalent bond

5. sp3 hybridization 4 atomic orbitals s + px + py + pz
Orbitals have two lobes (unsymmetrical)
Orbitals arrange in space with larger lobes away from one another (tetrahedral shape)
Each hybrid orbital holds 2e-
→ 4 equivalent hybrid orbitals
→ 4 sppp = 4 sp3

6. Formation of methane
The sp3 hybrid orbitals on C overlap with 1s orbitals on 4 H atoms to form four identical C-H bonds
Each C–H bond has a strength of 438 kJ/mol and length of 110 pm
Bond angle: each H–C–H is 109.5°, the tetrahedral angle.

7. Motivation for hybridization?
Better orbital overlap with larger lobe of sp3 hybrid orbital then with unhybridized p orbital
Stronger bond
Electron pairs farther apart in hybrid orbitals
Lower energy

8. Another example: ethane
C atoms bond to each other by overlap of an sp3 orbital from each
Three sp3 orbitals on each C overlap with H 1s orbitals
Form six C–H bonds
All bond angles of ethane are tetrahedral

9. Both methane and ethane have only single bonds Sigma (s) bonds
Electron density centered between nuclei
Most common type of bond
Pi (p) bonds
Electron density above and below nuclei
Associated with multiple bonds
Overlap between two p orbitals
C atoms are sp2 or sp hybridized

10. Bond rotation Single (s) bonds freely rotate
Multiple (p) bonds are rigid

11. sp2 hybridization

12. sp2 hybridization 4 atomic orbitals → s + px + py + pz →
Shape = trigonal planar (bond angle = 120º)
Remaining p orbital is perpendicular to the plane
3 equivalent hybrid orbitals + 1 unhybridized p orbital
3 spp + 1 p = 3 sp2 + 1 p

13. Formation of ethene (C2H4)
Two sp2-hybridized orbitals overlap to form a s bond
Two sp2 orbitals on each C overlap with H 1s orbitals
Form four C–H bonds
p orbitals overlap side-to-side to form a  bond
sp2–sp2 s bond and 2p–2p  bond result in sharing four electrons and formation of C-C double bond
Shorter and stronger than single bond in ethane

14. sp hybridization 4 atomic orbitals → s + px + py + pz →
Shape = linear (bond angle = 180º)
Remaining p orbitals are perpendicular on y-axis and z-axis
2 equivalent hybrid orbitals unhybridized p orbitals
2 sp + 2 p

15. Formation of acetylene (C2H2)
Two sp-hybridized orbitals overlap to form a s bond
One sp orbital on each C overlap with H 1s orbitals
Form two C–H bonds
p orbitals overlap side-to-side to form two  bonds
sp–sp s bond and two p–p  bonds result in sharing six electrons and formation of C-C triple bond
Shorter and stronger than double bond in ethylene

16. Summary of Hybridization
Hybridization of C
sp3
sp2 sp
Example
Methane, ethane
Ethene
Ethyne
# Groups bonded to C 4 3 2
Geometry
Tetrahedral
Trigonal planar
Linear
Bond angles
109.5
~120
~180
Types of bonds to C 4s
3s, 1p
2s, 2p
C-C bond length (pm)
154
133
120
C-C bond strength (kcal/mol) 90
146
200
C-H bond length (pm)
110
108
106

17. Hybridization of Heteroatoms
Same theory
Look at number of e- groups to determine hybridization
Lone pairs will occupy hybrid orbital
Ammonia:
N’s orbitals (sppp) hybridize to form four sp3 orbitals
One sp3 orbital is occupied by two nonbonding electrons, and three sp3 orbitals have one electron each, forming bonds to H
H–N–H bond angle is 107.3°
Water
The oxygen atom is sp3-hybridized
The H–O–H bond angle is 104.5°