1. Molecular shapes A simple matter of balls and sticks

2. Learning objectives  Describe underlying principles that govern theories of molecular shapes  Use Lewis dot diagrams to predict shapes of molecules using VSEPR

3. Valence shell electron pair repulsion  In order to understand properties like polarity, we need to predict molecular shapes  Lewis dot structure provides 2D sketch of the distribution of the valence electrons among bonds between atoms and lone pairs; it provides no information about the shape of the molecule

4. A hierarchy of models  VSEPR  Consider the problem in terms of electrostatic repulsion between groups of electrons (charge clouds, domains)  Valence bond theory  Acknowledges the role of orbitals in covalent bonding  Molecular orbital (MO) theory (the “real” thing)  Accommodates delocalization of electrons - explains optical and magnetic properties

5. Electron groups (clouds) minimize potential energy  Valence shell electron pair repulsion (VSEPR)  Identify all of the groups of charge: non-bonding pairs and bonds (multiples count as one)  Distribute them about the central atom to minimize potential energy (maximum separation of the groups)  This specifies the electronic geometry (also known as electron domain geometry or sometimes confusingly as molecular geometry)

6. Choices are limited  Groups (domains) of charge range from 2 – 6  Only one electronic geometry in each case  However, more than one molecular shape follows from electronic geometry depending on number of lone pairs  One surprise: the lone pairs occupy more space than the bonded atoms (with very few exceptions)  Manifested in bond angles (examples follow)  Molecular shape selection (particularly in trigonal bipyramid)

7. Two groups: linear  Except for BeH 2 (Be violates octet rule), all cases with two groups involve multiple bonds

8. Three groups: trigonal planar  Two possibilities for central atoms with complete octets:  Trigonal planar (H 2 CO)  Bent (SO 2 )  BCl 3 provides example of trigonal planar with three single bonds  B is satisfied with 6 electrons – violates octet rule

9. Four groups: tetrahedral  Three possibilities:  No lone pairs (CH 4 ) - tetrahedral  One lone pair (NH 3 ) – trigonal pyramid  Two lone pairs (H 2 O) – bent  Lone pairs need space: H-N-H angle 107°H-N-H angle 107° H-O-H angle 104.5°H-O-H angle 104.5° Tetrahedral angle 109.5°Tetrahedral angle 109.5°

10. Representations of the tetrahedron

11. Five groups of charge: trigonal bipyramid – most variations  Two different positions:  Three equatorial  Two axial  Equatorial positions are lower energy:  Lone pairs require occupy these locations preferentially

12. Five bonds, no lone pairs

13. Four bonds, one lone pair  Lone pair dictates geometry: equatorial position has lower energy than axial

14. Three bonds, two lone pairs  Both lone pairs occupy equatorial positions – lower energy than in axial

15. Two bonds, three lone pairs  The trend continues: all equatorial positions filled – lowest energy

16. Octahedron has six identical positions and high symmetry

17. No lone pairs  High symmetry

18. One lone pair  All positions are equally probable  Symmetry reduced

19. Two lone pairs  Minimum energy has axial symmetry, lone pairs lie along straight line

22. Molecules with multiple centers  A central atom is any atom with more than one atom bonded to it  Perform exercise individually for each atom  Electronic geometry and molecular shape will refer only to the atoms/lone pairs immediately attached to that atom

23. Taking it to the next level: acknowledging orbitals  VSEPR is quite successful in predicting molecular shapes based on the simplistic Lewis dot approach  But our understanding of the atom has the electrons occupying atomic orbitals  How do we reconcile the observed shapes of molecules with the atomic orbital picture of atoms