Molecular shapes
A simple matter of balls and sticks
Learning objectives
 Describe underlying principles that govern
theories of molecular shapes
 Use Lewis dot diagrams to predict shapes of
molecules using VSEPR
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
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
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)
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)
Two groups: linear
 Except for BeH2 (Be violates octet rule), all cases
with two groups involve multiple bonds
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
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°
Representations of the tetrahedron
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
Five bonds, no lone pairs
Four bonds, one lone pair
 Lone pair dictates geometry: equatorial position
has lower energy than axial
Three bonds, two lone pairs
 Both lone pairs occupy equatorial positions –
lower energy than in axial
Two bonds, three lone pairs
 The trend continues: all equatorial positions filled –
lowest energy
Octahedron has six identical
positions and high symmetry
No lone pairs
 High symmetry
One lone pair
 All positions are equally probable
 Symmetry reduced
Two lone pairs
 Minimum energy has axial symmetry, lone pairs lie
along straight line
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
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