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AP Chemistry
Chapter 9
Vocab (Ch 9)
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VSEPR- Valence Shell e- Pair Repulsion
bonding pair
non bonding pair – lone pair of electrons
electron domain – regions around a central atom
where e- are likely to be found.
molecular geometry- the arrangement of atoms in
space
electron domain geometry- the arrangement of edomains about the central atom of a molecule
The Molecular geometry is a derivation of the
Electron-Domain geometry
See Table 9.2 (page 309)
Electron Domains
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The e- in a multiple bond constitute a
single e- domain.
# of e- domains =
(# of atoms bonded to the central
atom) + (# of non bonding pairs on the
central atom)
Page 306
Molecular Shapes Website
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See VSEPR table handout for molecular
shapes
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http://www.molecules.org/VSEPR_table.html
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See B& L page 307-309
Effect of Non bonding e- and
multiple bonds on Bond Angles
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e- domains for non-bonding e- pairs exert
greater repulsive forces on adjacent edomains and thus tend to compress the bond
angles
e- domains for multiple bonds exert a greater
repulsive force on adjacent
e- domains than do single bonds.
Page 310
Molecules with Expanded
Valence Shells
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These shapes generally contain axial and
equatorial positions
See B&L pg. 312
Variations of the trigonal bipyramidal shape
show lone electron pairs in the equatorial
position
Variations of the octahedral shape show lone
electron pairs in the axial positions
Page 311
Molecules With More Than One
Central Atom
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You can use the VSEPR theory for
molecules with more than one central
atom, such as, CH3NH2.
Pages 313-314
Bond Polarity
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Bond Polarity is a measure of how equally the
e- in a bond are shared between the 2 atoms
of the bond.
Polarity is used when talking about covalently
bonded molecules.
If the molecule has only 2 different atoms,
such as, HF or CCl4 you can calculate the
electronegativity difference and determine
the type of covalent bond (polar or nonpolar).
Polarity and Bond Type
Electronegativity
Difference
Bonding Type
<0.5
Non-polar covalent
0.5 – 1.9
Polar covalent
> 2.0
ionic
Dipole Moment
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Dipole Moment – the measure of the amount
of charge separation in the molecule.
For a molecule with more than 2 atoms, the
dipole moment depends on both the polarities
of the individual bonds and the geometry of
the molecule.
The overall dipole moment of a polyatomic
molecule is the sum of its bond dipoles.
See B&L page 315 figure 9.9
Dipole Moment
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For each bond in the molecule, consider the
bond dipole (the dipole moment due only to
the 2 atoms in that bond)
The dipole “arrow” should point toward the
more electronegative atom in the bond
The overall dipole moment of a polyatomic
molecule is the sum of its bond dipoles.
(Consider the magnitude and direction of the
bond dipoles)
Different Theories
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VSEPR Theory (using Lewis Dot Structures)
Valence Bond Theory (using hybridization)
Molecular Orbital Theory (shows allowed states for ein molecules)
Go to the following web-site for a compare and
contrasting of the 3 different theories
http://www.chem.ufl.edu/~chm2040/Notes/Chapter_12/theory.
html
Hybridization
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An atom in a molecule may adopt a
different set of atomic orbitals (called
hybrid orbitals) than those it has in the
free state.
See B&L pages 319-322 for explanation
and diagrams of electron promotion
sp Hybrid Orbitals
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See section 9.5
pages 318-320
Look at the hybridization links on my
web page, they contain content,
pictures and animations.
sp2 and sp3 Hybrid Orbitals
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See section 9.5
pages 320-322
d Orbital Hybridization
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See section 9.5
pages 322
Multiple Bonds and
Hybridization
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See section 9.6
pages 324-326
Delocalized π Bonding
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See section 9.6
pages 327-330
Sigma Bonds ( σ )
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Sigma bonds occur when the e- density
is concentrated between the 2 nuclei.
These are single covalent bonds.
Sigma bonds can form from the overlap
of an s orbital with another s orbital, an
s orbital with a p orbital, or a p orbital
with a p orbital.
Pi Bonds ( π )
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Overlap of 2 “p” orbitals oriented
perpendicularly to the inter-nuclear axis
This overlap results in the sharing of
electrons.
The shared electron pair of a pi bond
occupies the space above and below
the line that represents where the two
atoms are joined together.
Multiple Bond
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A multiple bond consists of one sigma
bond and at least one pi bond.
A double bond consists of one sigma
bond and one pi bond.
A triple bond consists of one sigma
bond and two pi bonds.
A pi bond always accompanies a sigma
bond when forming double and triple
bonds.
Sigma and Pi
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Most of the bonding that you have seen
so far has bonding e- that are localized.
σ and π bonds are associated with
the 2 atoms that form the bond (and
NO other atoms)
Delocalized bonding can occur in
molecules that have π bonds and
more than one resonance structure.
Molecular Orbital Diagram
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Energy Level Diagram (Molecular Orbital
Diagram)
The H2 molecule is the easiest molecule
to plot on the molecular orbital diagram
Whenever 2 atomic orbitals overlap,
2 molecular orbitals form (one is a
bonding orbital and one is an antibonding orbital).
This is not on the AP exam
Paramagnetism
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Molecules with one or more unpaired
electrons are attracted into a magnetic
field
The more unpaired electrons in species,
the stronger the force of attraction
This behavior is called paramagnetism
Diamagnetism
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Substances with no unpaired electrons
are weakly repelled from a magnetic
field
This property is called diamagnetism
Diamagnetism is much weaker than
paramagnetism
Problems to Try
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Ch 9
# 5-13, 15, 23, 27, 31, 32, 34, 40, 43, 44 (a
and c), 63
AP Exam Problems to Try
1999 # 8
2000 # 7 (last section)
2002 # 6
2003 # 8
2004 # 7 & # 8
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2006 #7
2007 #6