MOLECULAR GEOMETRY
Molecular shapes
 The shape of a molecule is determined by bond angles. The
lines made by joining the nuclei in a molecule are the bond
angles.
 To predict molecular shape, we assume valence electrons
repel each other. A molecule will adopt a 3D geometry
which minimizes these repulsions.
 This the Valence Shell Electron Pair Repulsion(VSEPR)
model.
The VSEPR model
 A covalent bond forms between atoms when a pair of
electrons occupies the space between them. This is a
bonding pair and the space is an electron domain.
 A nonbonding pair defines a domain located on one atom
 NH3 has 3 bonding and one nonbonding pair
 The arrangement of electron domains about the central
atom of an ABn molecule is its electron domain geometry
 5 electron domain geometries
o Linear- 2 domains
o Trigonal planar- 3 domains
o Tetrahedral- 4 domains
o Trigonal bipyramidal- 5 domains
o Octahedral- 6 domains
 Molecular geometry is the arrangement of atoms in space
 Predict the geometry
o Draw the Lewis structure
o Count the number of electron pairs around the
central atom
o Arrange the pairs to minimize repulsions
o Describe the geometry in terms of bonded atoms
o Multiple bonds count as one domain
o Treat nonbonding pairs as a domain for domain
geometry
Effect of nonbonding electrons and multiple bonds
 A bonding pair of electrons is attracted by 2 nuclei. They do
not repel as much as a nonbonding pair which is attracted
by one nuclei.
 Electron domains for nonbonding electrons tend to
compress bond angles
 Consider 3 molecules with tetrahedral electron domain
geometry: CH4, NH3 , H2O
 The H-X-H bond angle decreases from 109.5 degrees to 107
degrees to104.5 degrees
 Multiple bonds cause similar compression
 There are 11 basic molecular shapes:
o 3 atoms(AB2)
 linear
 bent
 sp hybridization
o 4 atoms(AB3)
 Trigonal planar
 Trigonal pyramidal
 T-shaped
 sp2 hybridization
o 5 atoms(AB4)
 Tetrahedral
 Square planar
 Seesaw
 sp3 hybridization
o 6 atoms (AB5)
 Trigonal bipyramidal
 Square pyramidal
 dsp3 hybridization
o 7 atoms (AB6)
 Octahedral
 d2sp3 hybridization
Covalent bonding and orbital overlap
 Valence bond theory
o A covalent bond forms when the orbitals of 2 atoms
overlap
o 2 electrons(usually 1 from each atom) are in the orbital
overlap
o As 2 nuclei approach each other their orbitals overlap
o As overlap increases, bond strength increases and
potential energy decreases.
o At some distance minimum energy is reached, this is the
bond length
o As the atoms get closer their nuclei repel each other and
energy increases. Bond length is the distance at which
attractive forces and repulsive forces are balanced.
 Hybrid orbitals
o Consider the BeF2 molecule. It has a 1s22s2
configuration. There are no unpaired electrons for
bonding. Atomic orbitals are not adequate to describe
molecules
o The F-Be-F bond angle is 180 degrees. We know that one
electron from Be is shared with 1 electron from F.
o We could promote a 2s electron to a 2p orbital to get
unpaired electrons. Based on the geometry of p orbitals
this does not explain the geometry of the molecule
o The geometry can be explained if a 2s and a 2p orbital
mix to form a sp hybrid orbital. Since the orbitals are
equivalent, the minimum energy position would occur at
180 degrees.
Multiple bonds
 Single bonds( each atom shares one electron) are termed sigma
bonds(). The electron density lies on the inter- nuclear axis
 Pi() bonds : electron density lies above and below the axis. A
double bond has 1 sigma and 1 Pi bond. A triple bond has 1
sigma and 2 Pi bonds.
 The p orbitals involved in Pi bonds come from unhybridized
orbitals
The Hydrogen Molecule
 When 2 atomic orbitals overlap, 2 molecular orbitals form.
 Therefore, 1s(H) + 1s(H) must result in 2 MO’s for H2
 One has electron density between the nuclei( bonding MO) and
one has little electron density on opposite sides of the nuclei.
 Bond order= ½( bonding electrons- antibonding electrons).
 H2 has 2 bonding and 0 antibonding. It has a bond order of
1(single bond). Double bond= bond order of 2
 Consider He2 : it would have 2 bonding and 2 antibonding
electrons. It would have a zero bond order( it can not exist)
Magnetic Properties
 Two types of magnetic behavior
o Paramagnetism(unpaired electrons in the molecule)
results in a strong magnetic attraction
o Diamagnetism ( no unpaired electrons) resulting in weak
magnetic repulsion