#### Transcript AP Chapter 9 Molecular Shapes

```AP Chapter 9
Molecular Geometry and Bonding
Theories
Molecular Shapes
• The 3-dimensional shapes and sizes of molecules
are determined by their bond angles and bond
lengths.
• Molecules with a central atom surrounded by n
atoms B, denoted ABn, adopt many different
geometric shapes, depending on the value of n,
and on the atoms involved.
• The majority have one of 5 basic shapes: linear,
trigonal pyramidal, tetrahedral, trigonal
bipyramidal and octahedryl.
The 5 basic shapes for
ABn molecules
from the 5 basic shapes. This one starts with a tetrahedral.
The VSEPR Model
• VSEPR – valence shell electron-pair repulsion
model.
• Molecular geometries are based on the
repulsions between electron domains. (Like
charges repel.)
• Shared pairs (bonded pairs) and unshared
pairs (nonbonding pairs) create electron
domains around the central atom.
VSEPR
• Electron domains remain as far apart as
possible.
• Nonbonding pairs of electrons (unshared
pairs) exert more repulsions than bonding
pairs (shared pairs.)
• Electron domains from multiple bonds exert
slightly more repulsions than those from
single bonds.
VSEPR
• The arrangement of electrons around a central
atom is called electron-domain geometry.
• The arrangement of atoms is called the
molecular geometry.
In general, each non-bonding pair,
single bond or multiple bond produces
an electron domain around the central
atom.
Balloons tied
together at their
ends naturally
lowest-energy
arrangement.
Electron Domains
• The best arrangement of a given number of
electron domains is one that minimizes the
repulsions among them. This is why the
balloon analogy – they have similar preferred
geometries.
Predicting Shapes
1. Draw the Lewis structure of the molecule or
ion and count the total number of electron
domains around the central atom.
2. Determine the electron-domain geometry by
arranging the electron domains around the
central atom so repulsions are minimized.
(Table 9.1)
3. Use the arrangement of the bonded atoms to
determine the molecular geometry.
The Effect of Nonbonding Electrons
and Multiple Bonds on Bond Angles
The bond angles decrease as the number
of nonbonding electron pairs increase.
Nonbonding pairs of electrons
• Nonbonding pairs of electrons take up more
space than bonding pairs.
• The electron domains for nonbonding electron
pairs exert greater repulsive forces on
adjacent electron domains and tend to
compress the bond angles.
Electron domains for multiple bonds exert
electron domains than single bonds.
Molecular Shape and Molecular
Polarity
Covalent Bonding and Orbital Overlap
• Valence bond theory states that covalent
bonds are formed when atomic orbitals on
neighboring atoms overlap one another.
Hybrid Orbitals
• This involves mixing s, p and sometimes d
orbitals.
• Linear = sp
• Trigonal planar = sp2
• Tetrahedral = sp3
• Trigonal bipyramidal = sp3d
• Octahedral = sp3d2
Multiple Bonds
• Sigma bonds (σ) – the overlap of 2 s orbitals
• Pi bonds (π) – the sideways overlap of p
orbitals.
π bonds in Benzene rings and resonance structures
Resonance
structures
Hybridization Involving Triple Bonds
Sigma and Pi Bonds
• Double bonds consist of one δ bond and one π
bond
• A triple bond consists of one δ bond and two
π bonds
• π bonds must lie in the same plane, therefore,
the presence of π bonds makes the molecule
slightly rigid.
Molecular Orbitals
• In the Molecular Orbital theory, electrons exist
in allowed energy states called molecular
orbitals (MO).
• They can be spread among all the atoms in a
molecule, they have a specific amount of
energy and can hold a maximum number of 2
electrons.
Molecular Orbitals
• The combination of two atomic orbitals leads
to 2 MOs, one with low energy and one with
higher energy.
• The lower energy MO concentrates the charge
density in the region between the nuclei and
is called a bonding molecular orbital.
• The higher energy MO excludes electrons
from the region between the nuclei and is
called antibonding molecular orbitals.
Bond orders
• When the appropriate number or electrons are
placed into the MOs, bond orders can be
calculated, which is half the distance between
the number of electrons in bonding MOs and the
number of electrons in the antibonding MOs.
• A bond order of 1 corresponds to a single bond
and so on. They can be fractions!
Paramagnetism and Diamagnetism
• Paramagnetism – the attraction of a molecule
by a magnetic field due to unpaired electrons.
• Diamagnetism – in molecules where all the
electrons are paired, the molecules exhibit a
weak repulsion from a magnetic field.
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