# UNIT 6B Practice Test – Molecular Shapes, Molecular Polarity

```UNIT 6B Practice Test – Molecular Shapes, Molecular Polarity, Bonding Theory
Total
Hybridization Bonding NonDomains
Domains Bonding
eDomains
2
sp
2
0
Electron Geometry
Molecular Geometry
Linear
Linear (180&deg;)*
3
sp2
3
0
Trigonal Planar
Trigonal Planar*
3
sp2
2
1
Trigonal Planar
Bent (120&deg;)
4
sp3
4
0
Tetrahedral
Tetrahedral*
4
sp3
3
1
Tetrahedral
Trigonal Pyramidal
4
sp3
2
2
Tetrahedral
Bent (109&deg;)
5
sp3d
5
0
Trigonal Bipyramidal
Trigonal Bipyramidal*
5
sp3d
4
1
Trigonal Bipyramidal
See-Saw
5
sp3d
3
2
Trigonal Bipyramidal
T-Shape
5
sp3d
2
3
Trigonal Bipyramidal
Linear
6
sp3d2
6
0
Octahedral
Octahedral*
6
sp3d2
5
1
Octahedral
Square Pyramidal
6
sp3d2
4
2
Octahedral
Square Planar*
Polarity of a Molecule:
 Must contain a polar bond/dipole (difference in electronegativity ≥ .5)
 Vector sum of all polar bonds must not be 0 (Molecule cannot have symmetrical dipoles)
 In chart above, molecular geometries with (*) can be symmetrical
Valence Bond Theory:
 A chemical bonding theory that explains the bonding between two atoms is caused by the overlap of
half-filled atomic orbitals. The two atoms share each other's unpaired electron to form a filled orbital to
form a hybrid orbital and bond together.
Pi π Bonds:
 overlap of the orbitals occurs above and below a line of the two nuclei of the atoms (present in double
and triple bonds)
Sigma σ Bonds:
 overlap of the orbitals occurs on the line between the two nuclei of the atoms (single bonds)