1. Molecular Geometry and
Hybridization of Atomic Orbitals
2. Molecular Geometry
 Diatomic molecules are the easiest to
visualize in three dimensions
 Diatomic molecules are linear
3. Valence Shell Electron Pair Repulsion
• The ideal geometry of a molecule is
determined by the way the electron pairs
orient themselves in space
• The orientation of electron pairs arises from
• The electron pairs spread out so as to
 Frequently, we will describe two
geometries for each molecule.
1. Electronic geometry is determined by
the locations of regions of high electron
density around the central atom(s).
2. Molecular geometry determined by the
arrangement of atoms around the central
Electron pairs are not used in the
molecular geometry determination just the
positions of the atoms in the molecule are
5. The Valence Shell Electron Pair Repulsion
model predicts shapes.
1. e- pairs stay as far apart as possible to minimize
2. The shape of a molecule is governed by the
number of bonds and lone pairs present.
3. Treat a multiple bond like a single bond when
determining a shape. Each is a single e-group.
4. Lone pairs occupy more volume than bonds.
Predicting Molecular Shapes: VSEPR
6. Predicting Molecular Shapes
1. Draw Lewis structure
2. Determine the number of electron pairs
around the central atom. Count a multiple
bond as one pair.
3. Arrange electron pairs as shown in the
8. Basic shapes that minimize repulsions:
If the molecule contains:
• only bonding pairs – the angles shown are correct.
• lone pair/bond mixtures – the angles change a little.
 lone pair/lone pair repulsions are largest.
 lone pair/bond pair are intermediate in strength.
 bond/bond interactions are the smallest.
 Illustrate the geometry of the following
10. Molecular Geometry
1 Lone pair to lone pair is the strongest repulsion.
2 Lone pair to bonding pair is intermediate
3 Bonding pair to bonding pair is weakest
 Mnemonic for repulsion strengths
lp/lp > lp/bp > bp/bp
 Lone pair to lone pair repulsion is why bond
angles in water are less than 109.5o
12. Bond Angles and Lone Pairs
 Ammonia and water show smaller bond
angles than predicted from the ideal
 The lone pair is larger in volume than a bond
 There is a nucleus at only one end of the
bond so the electrons are free to spread out
over a larger area of space
13. The A-X-E Notation
 A denotes a central atom
 X denotes a terminal atom
 E denotes a lone pair
 O is central
 Two lone pairs
 Two hydrogen
14. Multiple Bonds
15. Molecular Geometry Summary with Lone
16. The steps in determining a molecular shape.
Count all e- groups around central
Note lone pairs and double
Count bonding and
Valence Bond (VB) Theory
 Covalent bonds are formed by the overlap of
 Atomic orbitals on the central atom can mix and
exchange their character with other atoms in a
 Process is called hybridization.
 Hybrids are common:
1. Pink flowers
 Hybrid Orbitals have the same shapes as
predicted by VSEPR.
Valence Bond (VB) Theory
2 Linear sp
4 Tetrahedral sp3
6 Octahedral sp3d2
19. Valence Bond Theory
 Unpaired electrons from one atom pair
with unpaired electrons from another atom
and give rise to chemical bonds
 Simple extension of orbital diagrams
20. Multiple Bonds
21. Hybrid Orbitals
 Hybridization of the s and p orbitals on carbon.
 The four sp3 hybrid orbitals have equal energy.
 The four valence electrons are distributed evenly
across the sp3 hybrid orbitals.
 The angle between the sp3 hybrid orbitals is
22. Hybrid Orbitals
The number of hybrid orbitals obtained equals the
number of atomic orbitals mixed.
The type of hybrid orbitals obtained varies with the
types of atomic orbitals mixed.
sp sp2 sp3 sp3d sp3d2
Types of Hybrid Orbitals
23. Figure 11.2 The sp hybrid orbitals in gaseous BeCl2.
orbital box diagrams
24. Figure 11.3 The sp2 hybrid orbitals in BF3.
25. Figure 11.4 The sp3 hybrid orbitals in CH4.
26. Figure 11.5 The sp3 hybrid orbitals in NH3.
27. Figure 11.5
The sp3 hybrid orbitals in H2O.