1. Molecular Shape VSEPR (Valence Shell Electron Pair Repulsion) Theory

2. VSEPR Theory VSEPR (Valence Shell Electron Pair Repulsion) Theory allows us to predict the shape of molecules. Each molecule has pairs of electrons- The bonding pairs form covalent bonds. The bonding pairs form covalent bonds. The nonbonding pairs occupy the same orbital. The nonbonding pairs occupy the same orbital. Each of these pairs can be thought of as an electron cloud- a cloud of negative charge.

3. Like Charges Repel These clouds of electrons, areas of negative charge repel each other in such a way as to get as far apart from each other as possible. When there are only two electron clouds, the farthest they can get away from each other is 180 o, a straight line.

4. Three Clouds With three electron clouds, the farthest the electron clouds can get away from each other is 120 o, the corners of an equilateral triangle. This shape is known as trigonal planer.

5. Four clouds When four clouds are attached to a central atom, the farthest they can get away from each other is the four corners of a tetrahedron. The tetrahedral angle is 109.5 o.

6. Non-bonding clouds also repel both bonding and non-bonding clouds. NH 3, ammonia, has a non-bonding pair of electrons. The central atom, N, is surrounded by 3 bonding and 1 non-bonding pair. These four clouds form roughly a tetrahedron, but because non-bonding pairs “spread out” more than bonding pairs, the angle between atoms is 107 o instead of 109.5 o. The atoms form a pyramidal shape.

7. H 2 O, water, has 2 bonding and 2 non-bonding pairs of electrons. These four clouds form roughly a tetrahedron, but because non-bonding pairs “spread out” more than bonding pairs, the angle between atoms is 105 o instead of 109.5 o. The atoms form a bent shape.

8. Summary 2 electron clouds: linear shape 3 electron clouds: trigonal planar shape 4 bonding clouds: tetrahedral shape 4 clouds- 3 bonding and one non-bonding: pyramidal shape 4 clouds- 2 bonding and 2 non-bonding: bent shape bent shape

9. The shape of the molecule affects the polarity of a molecule. Of course, if a molecule has only non- polar bonds, the molecule will be non- polar. If a molecule has polar bonds, the molecule itself may be polar or non-polar depending on its shape. If the charges are arranged in such a way as to cancel each other out, a molecule with polar bonds may be non-polar.

10. Note that all these molecules have polar bonds. But 2 molecules are nonpolar because their symmetry causes the charges to cancel out: CO2 (linear shape) and CCl4 (tetrahedral shape). Copyright by the Glencoe Division of Macmillan/McGraw-Hill School Publishing Company