1. Lesson 3.4 - Valence Shell Electron Pair Repulsion Theory (VSEPR)
Electron pairs repel as much as possible in 3-dimensional space.
Theory used to predict shapes of molecules.

3. To determine the shape of a molecule, we distinguish between lone pairs (or non-bonding pairs, those not in a bond) of electrons and bonding pairs (those found between two atoms).
The electrons adopt an arrangement in space to minimize e--e- repulsion.

5. Examples – Draw the Lewis structures, and then determine the orbital geometry of each:
H2S
CO2
PCl3
CH4
SO2

6. Covalent Bond – An attraction between two atoms caused by the sharing of a pair of electrons between two atoms.
Polar Covalent – A covalent bond in which electrons are shared unequally.
Nonpolar covalent – electrons are shared equally. SYMMETRICAL arrangement of valence electrons.

7. > 1.5 Ionic 0.5-1.5 Polar covalent < 0.5 Nonpolar covalent
Electronegativity – The tendency of an atom in a bond to attract shared electrons to itself.
Look at electronegativity difference to determine bond type:
EN Difference
Bond Type
> 1.5
Ionic
Polar covalent
< 0.5
Nonpolar covalent

8. Partially Positive
Partially Negative

9. Polar Molecule – A molecule in which valence electrons (bonds and unshared pairs of electrons) are not equally distributed (asymmetrical)
To be polar, a molecule must:
1. contain polar covalent bonds
2. be asymmetrical

10. Examples – Is each molecule polar or nonpolar?
CO2
H2S
CCl4
NBr3