Molecular shapes
A lot 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
Counting sheep clouds
 Valence shell electron pair repulsion (VSEPR)
 Identify all groups of charge on central atom only:
• Non-bonding pairs count 1
• All bonded atoms count 1 (singles/doubles/triples)
 Distribute them about central atom to minimize potential
energy (equals maximum separation)
 This specifies electronic geometry (also known as
electron domain geometry or sometimes, confusingly,
as molecular geometry)
Electronic geometry and molecular
shape
 Electronic geometry includes all
atoms and lone pairs on central
atom
 H2O has tetrahedral electronic
geometry (2 +2 = 4)
 Molecular shape (geometry)
ignores lone pairs
 H2O is bent (2 atoms)
 Must know electronic geometry
to obtain correct molecular
shape
 With no lone pairs: molecular
shape = electronic geometry
Blob counting: Choices are limited
 Groups (domains) of charge range from 2 – 6
 Only one electronic geometry for each number
 Note: more than one molecular shape follows from
electronic geometry depending on number of lone
pairs
 One surprise: 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 – the tricky one)
Two groups: linear
 Except for BeX2 (Be violates octet rule), all cases
with two groups involve multiple bonds
 Electronic geometry = molecular shape = linear
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 need more space
 Lone pairs require occupy equatorial sites 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