Chemistry for Engineering Students

1. Molecular Geometry and
Hybridization of Atomic Orbitals

2. Molecular Geometry
 Diatomic molecules are the easiest to
visualize in three dimensions
 HCl
 Cl2
 Diatomic molecules are linear

3. Valence Shell Electron Pair Repulsion
Theory
• 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
electron repulsions
• The electron pairs spread out so as to
minimize repulsion

4. 4
VSEPR Theory
 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
atom(s).
Electron pairs are not used in the
molecular geometry determination just the
positions of the atoms in the molecule are
used.

5. The Valence Shell Electron Pair Repulsion
model predicts shapes.
1. e- pairs stay as far apart as possible to minimize
repulsions.
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
next slide

7. Predicting Molecular Shapes:
Linear Triangular planar Tetrahedral
Triangular bipyramidal Octahedral

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.
linear triangular
planar
tetrahedral triangular
bipyramidal
octahedral

9. Examples
 Illustrate the geometry of the following
molecules:
1. BeH2
2. CH4
3. BF3
4. PCl5
5. SF6

10. Molecular Geometry

12. 12
VSEPR Theory
1 Lone pair to lone pair is the strongest repulsion.
2 Lone pair to bonding pair is intermediate
repulsion.
3 Bonding pair to bonding pair is weakest
repulsion.
 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
.

13. Bond Angles and Lone Pairs
 Ammonia and water show smaller bond
angles than predicted from the ideal
geometry
 The lone pair is larger in volume than a bond
pair
 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

14. The A-X-E Notation
 A denotes a central atom
 X denotes a terminal atom
 E denotes a lone pair
 Example
 Water
 H2O
 O is central
 Two lone pairs
 Two hydrogen
 AX2E2

15. Multiple Bonds

16. Molecular Geometry Summary with Lone
Pairs Included

19. The steps in determining a molecular shape.
Molecular
formula
Lewis
structure
Electron-group
arrangement
Bond
angles
Molecular
shape
(AXmEn)
Count all e- groups around central
atom (A)
Note lone pairs and double
bonds
Count bonding and
nonbonding e-
groups separately.
Step 1
Step 2
Step 3
Step 4

20. 20
Valence Bond (VB) Theory
 Covalent bonds are formed by the overlap of
atomic orbitals.
 Atomic orbitals on the central atom can mix and
exchange their character with other atoms in a
molecule.
 Process is called hybridization.
 Hybrids are common:
1. Pink flowers
2. Mules
 Hybrid Orbitals have the same shapes as
predicted by VSEPR.

21. 21
Valence Bond (VB) Theory
Regions of
High Electron
Density
Electronic
Geometry
Hybridization
2 Linear sp
3 Trigonal
planar
sp2
4 Tetrahedral sp3
5 Trigonal
bipyramidal
sp3d
6 Octahedral sp3d2

22. 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

24. Multiple Bonds

26. 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
109.5o.

27. 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.
Key Points
sp sp2 sp3 sp3d sp3d2
Types of Hybrid Orbitals

28. Figure 11.2 The sp hybrid orbitals in gaseous BeCl2.
atomic
orbitals
hybrid
orbitals
orbital box diagrams

29. Figure 11.3 The sp2 hybrid orbitals in BF3.

30. Figure 11.4 The sp3 hybrid orbitals in CH4.

31. Figure 11.5 The sp3 hybrid orbitals in NH3.

32. Figure 11.5 continued The sp3 hybrid orbitals in H2O.

33. Figure 11.6 The sp3d hybrid orbitals in PCl5.

34. Figure 11.7 The sp3d2 hybrid orbitals in SF6.

35. Hybrid Orbitals and Geometry