Transcript Chapter 9 Molecular Geometry and Bonding Theories

Chemistry 100
Chapter 9
Molecular Geometry and Bonding
Theories
Molecular Geometry
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The three-dimensional arrangement of atoms
in a molecule  molecular geometry
Lewis structures can’t be used to predict
geometry
A very simple theory tells us that the
repulsion between electron pairs (both
bonding and non-bonding) helps account for
the arrangement of atoms in molecules
The VSEPR Model
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Electrons are negatively charged, they want to occupy
positions such that electron – electron interactions are
minimised as much as possible
Valence Shell Electron-Pair Repulsion Model
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treat double and triple bonds as single bonds
resonance structure - apply VSEPR to any of them
Formal charges are usually omitted
Molecules With More Than One
Central Atom
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We simply apply VSEPR to each ‘central atom’ in the
molecule.
• Carbon #1 – tetrahedral
• Carbon #2 – trigonal
planar
Dipole Moments
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The HF molecule has a bond dipole – a charge
separation due to the electronegativity difference
between F and H.
The shape of a molecule and the magnitude of the bond
dipole(s) can give the molecule an overall degree of
polarity  dipole moment.
+H-F
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Homonuclear diatomics  no dipole moment (O2, F2,
Cl2, etc)
Triatomic molecules (and greater). Must look at the net
effect of all the bond dipoles.
In molecules like CCl4 (tetrahedral) BF3 (trigonal
planar) all the individual bond dipoles cancel  no
resultant dipole moment.
Bond Dipoles in Molecules
More Bond Dipoles
Valence Bond Theory and
Hybridisation
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Valence bond theory
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description of the covalent bonding and structure in
molecules.
Electrons in a molecule occupy the atomic
orbitals of individual atoms.
The covalent bond results from the overlap of
the atomic orbitals on the individual atoms
The H2 Molecule
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In a hydrogen molecule, we observe a single
bond indicating the overlap of the 1s orbitals
on the individual atoms
The Bonding in H2
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Note that bond has cylindrical symmetry with respect
to the line joining the atomic centres description. This
is known as a  bond
H
H
Overlap Region
1s (H1) – 1s(H2)  bond
The Bonding in H2
H
H
The Cl2 Molecule
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In the chlorine molecule, we observe a single
bond indicating the overlap of the 3p orbitals
on the individual atoms.
Cl
Cl
Bonding description 3pz (Cl 1) – 3pz (Cl 2)
Is This a  Bond?
Cl
Cl
Hybrid Atomic Orbitals
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Look at the bonding picture in methane
(CH4).
Bonding and geometry in polyatomic
molecules may be explained in terms of the
formation of hybrid atomic orbitals
Bonds - overlap of the hybrid atomic
orbitals with the atoms. appropriate halffilled atomic orbital on the terminal
The CH4 Molecule
The Formation of the
3
sp
Hybrids
2
sp
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Hybridisation
What if we try to rationalise the bonding
picture in the BH3 (a trigonal planar
molecule)?
We mix 2 “pure” p orbitals and a “pure” s
orbital to form “hybrid” or mixed sp2
orbitals.
These three sp2 hybrid orbitals lie in the
same plane with an angle of 120 between
them.
A Trigonal Planar Molecule
H
H
Overlap regions
B
Overlap region
H
sp Hybridisation
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What if we try to rationalise the bonding
picture in the BeH2 species (a linear
molecule)?
We mix a single “pure” p orbital and a “pure”
s orbital to form two “hybrid” or mixed sp
orbitals
These sp hybrid orbitals have an angle of
180 between them.
A Linear Molecule
The BeH2 molecule
Overlap Regions
H
Be
H
Double Bonds
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Look at ethene C2H4.
Each central atom is an AB3 system, the
bonding picture must be consistent with
VSEPR theory.
The  Bond
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Additional feature
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an unhybridized p orbital on adjacent carbon
atoms.
Overlap the two parallel 2pz orbitals (a orbital is formed).
The C2H4 Molecule
The Bond Angles in C2H4
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Bond angles HCH = HCC  120. Note
that the  bond is perpendicular to the plane
containing the molecule.
We can rationalize the presence of any
double bond by assuming sp2 hybridization
exists on the central atoms!
Any double bond  one  bond and a  bond
The Triple Bond
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Look at acetylene (ethyne)
•The carbon atoms each have a triple bond
and a single bond.
The C2H2 Molecule
The  Bonds
• On the carbon central atom, we now have 2
sp hybrid orbitals and two unhybridised p
orbitals
• We can again overlap the 2py orbitals and the
2pz orbitals on the C central atoms (two pbonds are formed).
The Bond Angles in C2H2
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Bond angles HCH = HCC = 180. The 
bonds are again perpendicular to the plane
containing the molecule.
Triple bond  one  bond and two  bonds
Rationalise the presence of any triple
bond by assuming sp hybridization
exists on the central atoms!
3
sp d
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Hybridisation
How can we use the hybridisation concept to explain the
bonding picture PCl5.
There are five bonds between P and Cl (all  type
bonds).
5 sp3d orbitals  these orbitals overlap with the 3p
orbitals in Cl to form the 5  bonds with the required
VSEPR geometry  trigonal bipyramid.
Bond overlaps
[sp3d (P ) – 3pz (Cl) ] x 5
 type
3
2
sp d
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Hybridisation
Look at the SF6 molecule.
6 sp3d2 orbitals  these orbitals overlap with
the 2pz orbitals in F to form the 6  bonds
with the required VSEPR geometry 
octahedral.
Bond overlaps
[sp3d2 (S ) – 2pz (F) ] x 6
 type
Notes for Understanding
Hybridisation
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Applied to atoms in molecules only
Number hybrid orbitals = number of atomic orbitals used
to make them
Hybrid orbitals have different energies and shapes from
the atomic orbitals from which they were made.
Hybridisation requires energy for the promotion of the
electron and the mixing of the orbitals  energy is offset
by bond formation.
Delocalised Bonding
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In almost all the cases where we described the bonding
n the molecule, the bonding electrons have been totally
associated with the two atoms that form the bond  they
are localised.
What about the bonding situation in benzene, the nitrate
ion, the carbonate ion?
In benzene, the C-C  bonds are formed from the sp2
hybrid orbitals. The unhybridised 2pz orbital on C
overlaps with another 2pz orbital on the adjacent C atom.
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Three  bonds are formed. These  bonds
extend over the whole molecule (i.e. the 
bonds are delocalised).
The  electrons are free to move around the
benzene ring.
Any species where we had several
resonance structures, we would have
delocalisation of the -electrons.
Delocalised Electrons in Molecules
Molecular Orbital (M.O.) Theory
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Valence bond and the concept of the hybridisation of
atomic orbitals does not account for a number of
fundamental observations of chemistry.
To reconcile these and other differences, we turn to
molecular orbital theory (MO theory). In MO theory,
covalent bonding is described in terms of molecular
orbitals, i.e., the combination of atomic orbitals that
results in an orbital associated with the whole molecule.
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Recall the wave properties of electrons.
constructive interference  the two e- waves interact
favourably; loosely analogous to a build-up of edensity between the two atomic centres.
destructive interference  unfavourable interaction of
e- waves; analogous to the decrease of e- density
between two atomic centres.
Constructive and Destructive
Interference
Constructive
+
Destructive
+
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ybonding = C1 ls (H 1) + C2 ls (H 2)
yanti = C1 ls (H 1) - C2 ls (H 2)
Bonding Orbital  a centro-symmetric orbital (i.e.
symmetric about the line of symmetry of the bonding
atoms).
Bonding M’s have lower energy and greater stability than
the AO’s from which it was formed.
Electron density is concentrated in the region
immediately between the bonding nuclei.
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Anti-bonding orbital  a node (0 electron density)
between the two nuclei.
In an anti-bonding MO, we have higher energy and less
stability than the atomic orbitals from which it was
formed.
As with valance bond theory (hybridisation)
2 AO’s  2 MO’s
Bonding and Anti-Bonding M.O.’s
from 1s atomic Orbitals
* 1s
1s
1s
Energy
1s
The MO’s in the H2 Atom
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The situation for two 2s orbitals is the same! The
situation for two 3s orbital is the same.
Let’s look at the following series of molecules
H2, He2+, He2
bond order = ½ {bonding - anti-bonding e-‘s}.
Higher bond order  greater bond stability.