1. Molecular Shapes It mean the 3-D arrangement of atoms in a molecule
Lewis dot structures show how atoms are bonded together (2-D), but they often do not illustrate the true shape of a molecule (3-D).
How to determine the shape of a molecule?
 Valence Shell Electron Pair Repulsion Theory
(VSEPR)
 VSEPR determines if a molecule is polar!

2. VSEPR Theory
Pairs of electrons around a central atom want to be as far away from each other as possible. Main idea: Bonding and Non-bonding e- pairs are positioned as far as possible so as to minimize repulsions between electron pairs. For VSEPR, treat double and triple bonds like single bonds. e- pairs: Bonding and Non-bonding pairs (lone pairs) around the central atom, don’t include central atom itself

3. Steps for Predicting Molecular Structures using VSEPR Model
Draw the Lewis Structure for the molecule.
Count the electron pairs and arrange them in the way that minimize repulsions (put the pairs as far apart as possible).
Determine the positions of the atoms from the way the electrons pairs are shared.
Determine the name of the molecular structure from the positions of the atoms.
Must Show 3-D shape by using and

4. Be Careful When Drawing Molecular Shapes
No Charge and No [ ] around the shape
Only Draw the Lone Central Atom
 Don’t draw the lone pairs around other atoms
Bend when having lone central atom
When Bending, Lone Pairs stay on the top, Bonds stay on the bottom.
More Lone Central Atom  Bend More

5. Lone Pairs around Central Atoms
Lone pair: Valence e- pair that belongs to only 1 atom / is not involved in bonding
Linear, Trigonal Planar, and Tetrahedral have NO lone pairs around the central atoms
 Just look at how many atoms around the central atom
That is how many e- pairs

6. Lone Pairs around Central Atoms
What about if there are lone pairs around the central atom?
REMEMBER:
Lone pairs also take up space!!!
Lone pairs repel each other
Lone pairs repel bonded pairs
Use as Lone central atom instead
of electron pairs.

7. 5 Molecular Shapes
A) Linear (2 kinds) B) Trigonal Planar C) Tetrahedral D) Trigonal Pyramid E) Bent (2 kinds)

8. A) Linear Linear 2 Linear 1
If two atoms bond to a central atom, like in CO2, the molecule will form in a straight line (Linear).  2 e- pairs, 3 atoms, angle = 180 degrees
Linear 1
Linear 2

9. B) Trigonal Planar
If three atoms bond to a central atom, with no lone pairs, like in BF3 , the molecule forms a Trigonal Planar shape.  3 e- pairs, 4 atoms, angle = 120 degrees

10. Molecular Shapes
C) Tetrahedral
If four atoms bond to a central atom, like in CH4, the molecule takes a tetrahedral shape.  4 e- pairs, 5 atoms, angle = degrees

11. D) Trigonal Pyramidal
If one lone pair of electrons and three atoms bond to a central atom, like in CH4, the molecule takes on a Trigonal Pyramidal shape.  4 e- pairs (1 lone pair + 3 bonded pairs), 4 atoms, angle = degrees Remember, lone pairs take up space!
107.2

12. E) Bent
Bent 1 If one lone pairs of electrons and two atoms bond to central atom, like in SO2 3 e- pairs (1 lone pair + 2 bonded pairs), 3 atoms, angle = – 120 degrees
Bent 2 If two lone pairs of electrons and two atoms bond to central atom, like in H2S 4 e- pairs (2 lone pairs + 2 bonded pairs), 3 atoms, angle = 105 degrees