Transcript Chapter 9 Molecular Geometry and Bonding Theories

Chemistry 100 Chapter 9
Molecular Geometry and Bonding
Theories
Molecular Geometry

The three-dimensional arrangement of
atoms in a molecule  molecular geometry

Lewis structures can’t be used to predict
geometry

Repulsion between electron domains (both
bonding and non-bonding) helps account for
the arrangement of atoms in molecules
The VSEPR Model

Electrons are negatively charged, they want to
occupy positions such that electron
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Valence Shell Electron-Pair Repulsion Model
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Electron interactions are minimised as much as
possible
treat double and triple bonds as single domains
resonance structure - apply VSEPR to any of them
Formal charges are usually omitted
Molecules With More Than One
Central Atom

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 Bonding in H2

Hydrogen molecule

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a single bond indicating the overlap of the 1s
orbitals on the individual atoms
cylindrical symmetry with respect to the line
joining the atomic centres, i.e., a  bond
H
H
Overlap Region
1s (H1) – 1s(H2)  bond
The Bonding in H2
H
H
The Cl2 Molecule

In the chlorine molecule, we observe a
single bond indicating the overlap of the
3p orbitals on the individual atoms.
Cl
Bonding description
Cl
3pz (Cl 1) – 3pz (Cl 2)
Is This a  Bond?
Cl
Cl
Hybrid Atomic Orbitals

Look at the bonding picture in methane
(CH4).
Bonding and geometry in polyatomic systems
may be explained in terms of the formation of
hybrid atomic orbitals
Bonds - overlap of the hybrid atomic orbitals
with the atoms. appropriate half-filled atomic
orbital on the terminal
The CH4 Molecule
The Formation of the sp3
Hybrids
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We mix 3 “pure” p orbitals and a “pure”
s orbital to form a “hybrid” or mixed
orbital sp3.
This is how we can rationalise the
geometry of the bonds around the C
central atom.
How do we form the hybrid orbitals?
The Formation of the sp3
Hybrids
Bonding Picture in CH4
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In CH4, the carbon sp3 orbitals (4 of
them) overlap with the s orbitals on the
H atoms to form the C-H bond
Bond overlaps
[sp3 (C) – 1s (H) ] x 4
 type
sp2 Hybridisation
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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).
Bond overlaps in C2H4

There are three
different types of
bonds
[sp2 (C ) – 1s (H) ] x 4
 type
[sp2 (C 1 ) – sp2 (C 2 ) ]
 type
[2pz (C 1 ) – 2pz (C 2 ) ]
 type
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 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!
Bond Overlaps in C2H2
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There are again three different types of
bonds
[sp (C ) – 1s (H) ] x 2  type
[sp (C 1 ) – sp (C 2 ) ]  type
[2py (C 1 ) – 2py (C 2 ) ]
 type
[2pz (C 1 ) – 2pz (C 2 ) ]
 type
sp3d Hybridisation
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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
sp3d2 Hybridisation
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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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Most cases

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bonding electrons have been totally
associated with the two atoms that form the
bond  they are localized.
Benzene
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The C-C  bonds are formed from the sp2
hybrid orbitals.
The unhybridized 2pz orbital on C overlaps
with another 2pz orbital on the adjacent C
atom.
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Three delocalized  bonds are formed.
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 bonds extend over the whole molecule.
The  electrons are free to move around
the benzene ring.
Several resonance structures, we
would have delocalization 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.
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.

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
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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
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.
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