Molecular shapes

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 BeH2 (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 (H2CO) • Bent (SO2) • BCl3 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 (CH4) - tetrahedral • One lone pair (NH3) – trigonal pyramid • Two lone pairs (H2O) – bent • Lone pairs need space: • H-N-H angle 107° • H-O-H angle 104.5° • 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