A. Octet Rule n Remember… Most atoms form bonds in order to have 8 valence electrons.
Hydrogen 2 valence e - Groups 1,2,3 get 2,4,6 valence e - Expanded octet more than 8 valence e - (e.g. S, P, Xe) Radicals odd # of valence e - n Exceptions: A. Octet Rule
B. Drawing Lewis Diagrams n Find total # of valence e -. n Arrange atoms - singular atom is usually in the middle. n Form bonds between atoms (2 e - ). n Distribute remaining e - to give each atom an octet (recall exceptions). n If there aren’t enough e - to go around, form double or triple bonds.
D. Resonance Structures n Molecules that can’t be correctly represented by a single Lewis diagram. n Actual structure is an average of all the possibilities. n Show possible structures separated by a double-headed arrow.
D. Resonance Structures O O S O O O S O O O S O n SO 3
IIIIII II. Molecular Geometry Molecular Structure
A. VSEPR Theory n Valence Shell Electron Pair Repulsion Theory n Electron pairs orient themselves in order to minimize repulsive forces.
A. VSEPR Theory n Types of e - Pairs Bonding pairs - form bonds Lone pairs - nonbonding e - Lone pairs repel more strongly than bonding pairs!!!
A. VSEPR Theory n Lone pairs reduce the bond angle between atoms. Bond Angle
n Draw the Lewis Diagram. n Tally up e - pairs on central atom. double/triple bonds = ONE pair n Shape is determined by the # of bonding pairs and lone pairs. Know the 8 common shapes & their bond angles! B. Determining Molecular Shape
C. Common Molecular Shapes 2 total 2 bond 0 lone LINEAR 180° BeH 2
3 total 3 bond 0 lone TRIGONAL PLANAR 120° BF 3 C. Common Molecular Shapes