Download Chapter 10

11/30/2011
Chapter 10
Molecular Geometry
VSEPR Model: Geometries
Valence Shell Electron Pair Repulsion Theory
• Electron pairs repel and get as far apart as possible
• Example: Water
– Four electron pairs
– Farthest apart is a tetrahedron
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VSEPR Model: Geometries
• Experimentally, we only observe
the atoms and bonds between
them.
• Water has a bent geometry
VSEPR Model: Geometries
• Start with the Lewis structure
• Count number of repelling groups around an atom
– Bonds (single, double, triple)
– Lone pairs
• # repelling groups determines their arrangement
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VSEPR Model: Geometries
# Repelling Groups
Arrangement of
Groups
Approx. Bond Angles
2
Linear
180o
3
Trigonal
120o
4
Tetrahedral
109o
5
Trigonal Bipyramidal
120o, 90o
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Octahedral
90o
You should know the arrangement
and the approximate bond angles.
VSEPR Model: Geometries
• What is the arrangement of groups and the
geometry of each molecule below?
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VSEPR Model: Geometries
• XeF4
• Arrangement: octahedral
• Geometry: square planar
VSEPR Model: Geometries
• Acetone
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VSEPR Model: Geometries
• Ethanol
VSEPR Model: Bond Angles
• We can make better estimations of bond angles
– Some groups repel better than others
• Repelling power:
– Lone pair > triple bond > double bond > single bond
• Compare CH4, NH3, and H2O
– Arrangement of groups is tetrahedral in all cases
– Bond angles differ slightly
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VSEPR Model: Bond Angles
Double bonds have more repelling
power than single bonds.
Lone pairs are even
better “repellers”
VSEPR Model: Bond Angles
• Draw the Lewis structure of each molecule below.
• Draw a picture of each molecule, showing its
geometry and predicting approximate bond angles
• CH2CH2
• O3
• IF3
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Valence Bond Theory
• Quantum mechanical theory of bond formation
• Extension of Lewis theory
– Lewis: Atoms share a pair of electrons to form bond
– VB: Orbitals of atoms overlap to form bonds
• Only orbitals containing unpaired valence electrons
• Example
– HF
– Unpaired valence electrons: H = 1s
F = 2p
Valence Bond Theory
• Advantage of VB Theory over Lewis
– Explains why covalent bond is formed
– Increased electron density between positive nuclei
– Nuclei are attracted to electrons between them
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Sigma Bonds
• These bonds are classified as sigma bonds
• Rotational symmetry about internuclear axis
Hybrid Orbitals
• Our model of bond generation
–
–
–
–
2 electrons per bond
Only unpaired electrons involved
All atoms get full octet of electrons
Bonds formed by overlap of atomic orbitals
• Seems to conflict with carbon
– 1s22s22p2
– If we use only unpaired e-, how many bonds predicted?
– How many bonds do we know C always forms?
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Hybrid Orbitals
• Carbon always forms 4 bonds
• Explanation
– Atom forms new orbitals in process of forming bonds
1s
2s
2p
sp3
1s
– New orbitals are called hybrid orbitals
– Orientation of hybrid orbitals is same
as VSEPR group arrangements
Hybrid Orbitals
• Assume that hybrid orbitals always are formed
before bonding
• Can predict hybridization of central atom from
arrangement of e- pairs
Arrangement of e- pairs
Hybridization
2
sp
3
sp2
4
sp3
5
sp3d
6
sp3d2
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Hybrid Orbitals
• Example: CH4 (methane)
• Lewis structure:
• Four repelling groups
– Tetrahedral geometry
– sp3 hybridization
– Four hybrid orbitals
• Four sigma bonds
Hybrid Orbitals
• Predict hybridization of central atoms in following
molecules:
• O
in H2O
• Two C’s in
• S
CH3CHO
in SF6
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Multiple Bonds
• Example: Ethylene
• Each C has 3 repelling groups
– sp2 hybrid
1s
sp2
2p
• sp2 hybrids form σ-bonds with
H atoms and other C atom
Multiple Bonds
• What about the remaining unpaired e- in the 2p
orbital on each C atom?
• Parallel p-orbitals overlap and form a π-bond
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Multiple Bonds
• Single bond:
• Double bond:
• Triple bond:
one σ
one σ and one π
one σ and two π
• How many σ and π bonds in CH2=C=CH2?
• Check blog. I will add a question about this molecule.
Sigma and Pi Bonds
• Electron density in σ-bonds is concentrated between
nuclei
• Electron density in π-bonds is above and below
nuclei
• Electrons in a π-bond are less tightly held than those
in a σ-bond
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Sigma and Pi Bonds
• A π-bond is weaker than a σ-bond
– A double bond is stronger than a single bond, but not
twice as strong
• A π-bond is more reactive than a σ-bond
– Example: Bromination of 2-butene
Sigma and Pi Bonds
• Alkanes (CnH2n+2)
– Example: n-Butane: C4H10
– All bonds are sigma: strong and not
readily broken
– Alkanes are not very reactive
– Used as solvents and fuels
• Reactivity of organic compounds mainly dependent
on multiple bonds and functional groups
CH3CH2-OH
CH3-NH2
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Dipole Moments
• Polar bonds have unequal sharing of electrons
• The dipole moment is a measure of the polarity
– High dipole moment = high polarity
Dipole Moments
• Not all bonds are polar
• Identical atoms: F – F
– Completely nonpolar
• Atoms with similar electronegativities: C – H
– ΔEN ≤ 0.4
–
Dipole moment is very small
Bond is essentially nonpolar
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Molecular Dipole Moment
• Molecules have dipole moments
– Resultant of dipole moments of bonds
• Example: water
–
O–H bond: EN difference = 3.5 – 2.1 = 1.4 Polar
– Two polar bonds point in same direction
• Reinforce each other
– Water molecule is polar (has a significant
dipole moment)
Molecular Dipole Moments
• For a molecule to be polar
– It must have polar bonds
– Its geometry must not balance
• If all bonds are nonpolar or essentially nonpolar, then
the molecule is nonpolar.
– CH3CH2CH3
• C-C bond: Completely nonpolar
• C-H bond: ΔEN = 2.5 – 2.1 = 0.4 Essentially nonpolar
• The molecule is nonpolar
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Molecular Dipole Moment
• Polar bonds do not necessarily mean a molecule is
polar
– Polar bonds can balance each other
• Example: CO2
• The dipole moment of CO2 is zero
Molecular Dipole Moment
• If all bonds in an arrangement (linear, trigonal,
tetrahedral, etc) are identical, then the molecule
will be nonpolar
• Example: CF4
• C-F electronegativity difference
–
ΔEN = 4.0 – 2.5 = 1.5
Polar bond
• Identical bonds in all positions
– Dipole moments of bonds will balance
– Molecule has zero dipole moment
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Molecular Dipole Moment
• Rank these isomers in order of increasing dipole
moment
Molecular Dipole Moment
• Polar bonds are local properties of a molecule
• Overall polarity of molecule depends on relative
significance of polar and nonpolar groups
• Which of these molecules would be more polar?
– CH3CH2CH2CH2NH2
– CH3CH2NH2
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End of material for Exam 4
Molecular Orbital Theory
• Molecular orbital theory
– Looks at molecule as a whole
– All orbitals of atoms in molecule combine to form
molecular orbitals
• MO theory problems from text
– 52, 54, 55, 56, 58, 63, 64, 65, 68, 107
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Interference of Waves
• Constructive
– In phase
• Destructive
– Out of phase
• Noninteracting
– Physically separated
Molecular Orbital Theory
• Atomic orbitals combine to form molecular orbitals
• Three ways they can combine
– Antibonding: destructive interference
• Destabilizes molecule
– Nonbonding: non-interacting
• Neither stabilizes or destabilizes molecules
– Bonding: constructive interference
• Stabilizes molecule
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Molecular Orbital Theory
• Atomic orbitals combine to form molecular orbitals
• # molecular orbitals = # atomic orbitals
• 2 atomic orbitals => 2 molecular orbitals
– One bonding, one antibonding
• 3 atomic orbitals => 3 molecular orbitals
– One bonding, one nonbonding, one antibonding
MO Theory: H2 Molecule
• H2 molecule
– Combine two 1s atomic orbitals (one from each H atom)
– Get one bonding and one antibonding molecular orbital
• Bonding = σ; antibonding = σ*
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MO Theory: H2 Molecule
• Feed electrons from atomic orbitals into mo’s
– Each H atom: 1 e– H2 molecule: 2 e-
MO Theory: Bond Order
• Bond order: indicator of bond strength in molecule
• Bond order = ½(# bonding e- - # antibonding e-)
–
–
–
–
BO = 0
BO = 1
BO = 2
Etc.
Molecule not stable
Single bond
Double bond
• What is bond order for
H2 molecule?
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MO Theory: He2
• Does MO theory predict that He2 will be stable?
• What about He2+?
MO Theory: Homonuclear Diatomics
• There are higher MO energy levels
•
Li2, B2, C2, N2
O2, F2
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MO Theory: Homonuclear Diatomics
• What is the bond order of the carbide ion, C22- ?
– Use Li2, B2, C2, N2 energy level diagram
– Ten electrons
• Four from each carbon, two extra from -2 charge
MO Theory: Homonuclear Diatomics
• Predict the magnetism of N2 and O2.
– Dia- or paramagnetic?
• http://www.youtube.com/watch?v=yJs5ENtilIo
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MO Theory: Homonuclear Diatomics
• Make predictions about O2 molecule and ions
• Which has the strongest bond: O2,O2+, or O2-?
Why Use MO Theory?
• MO theory is used most of the time to calculate
properties of large molecules
– Predicts energy levels, so can be compared with electronic
spectroscopy experiments
– Better at predicting magnetic properties
– Better method of describing molecules for which electrons
are shared by more than two atoms
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MO Theory: Ozone
• O3: two resonance hybrids
• Each oxygen is sp2 hybridized
– Leaves one extra p-orbital on each oxygen
• Focus on these p-orbitals
MO Theory: Ozone
• Three p-orbitals
• Combine to form three molecular orbitals
– Antibonding
– Nonbonding
– Bonding
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MO Theory: Ozone
• Bonding molecular orbital is most important
• This is a delocalized molecular orbital
– Spread over three atoms
– Contains 2 electrons
Delocalized Molecular Orbitals
• Compounds that must be described by resonance
• Example: carbonate ion
• The lowest bonding molecular orbital usually is
delocalized
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Delocalized Molecular Orbitals
• Benzene, C6H6, is most important example
• Lowest energy bonding molecular orbital is
doughnut-shaped, above and below ring
Most common symbol for
Benzene is below.
Molecules with delocalized molecular orbitals generally are
more stable than molecules with localized double bonds
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