Some Material Copyright PGCC CHM 101 Sinex
Some Graphics from Nelson Chemistry 12 Textbook
Atomic Orbitals Don’t Work?
To explain molecular geometry:
• In methane, CH4 , the shape is
• The valence electrons of carbon should
be two in s, and two in p.
• The p orbitals would have to be at right
• The atomic orbitals change when making
Valence Bond Theory
• Atomic orbitals overlap when a bond is
formed between atoms
• A new set of orbitals is formed
• Each orbital contains a pair of electrons
• These are called hybrid orbitals
• How do they look? See p234 Table 1
• p234 Table 1
Ex: sp3 hybridization
• The s and p orbitals blend and end up
with the tetrahedral geometry.
• one s orbital combines with three p
• sp3 hybridization has tetrahedral
In Terms of Energy
• One s orbital and three p orbitals
• Each orbital holds 2 electrons
Animation of hybridization
Visualization of orbitals
Double and Triple Bonds
• Double bond acts as one pair.
• Geometry - trigonal planar
• Have to end up with three blended
• Use one s and two p orbitals to make sp2
• Leaves one p orbital perpendicular.
Where is the P orbital?
• The overlap of orbitals (end-to-end)
makes a sigma bond ( bond)
hybrid orbitals – sp, sp2, or sp3
formation of bond
remaining p orbitals form sp or sp2
bond hinders rotation Planar molecule (each carbon
about the carbon-to- is trigonal planar) with cloud
carbon bond above and below the plane
The overlap of orbitals
pi bond ( bond)
• End up with two lobes 180º apart.
• p orbitals are at right angles
• Makes room for two bonds and
• A triple bond or two double bonds.
Bond Formation From
• Single bond end-to-end
- sigma bond
• Double bond side-to-side
– sigma bond + pi bond
• Triple bond
– sigma bond + pi bond + pi bond
How are the two pi bonds in the triple bond oriented?
How about the electron density
around a C-C bond?
C2H6 C 2H 2
single bond double bond triple bond