Download Chapter 10 Molecular Geometry and Chemical Bonding Theory

VSEPR Theory
Valence Shell Electron Pair Repulsion
Chapter 10
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Molecular Geometry
and
Chemical Bonding Theory
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Different geometrical arrangements of
regions of electron density around a
molecule’s central atom would result;
dependent on the number of electron pairs.
VSEPR Theory and Molecular Geometries
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Imagine 2, 3, 4, 5, or 6 balloons tied together:
If we could see electron pairs, what would
this look like?
What 3–D shape would result?
VSEPR Theory and Molecular Geometries:
The Balloon Model
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In a polyatomic covalent species, electron pairs
arrange themselves in such a way as to
minimize electron – electron repulsion.
Now consider that each balloon represents the
space occupied by an electron pair (or region of
electron density; electron domain).
What is the geometrical arrangement of the
electron domains around a central atom?
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What shapes result?
What angles are created between the balloons by
their natural tendency to stay as far away from
each other as possible?
Electron Domain Geometry
Trigonal Bipyramidal Geometry
2 types of positions
around the central atom:
A – axial positions
E – equitorial positions
note the angles:
∠ A to E: 90°
∠ E to E: 120°
Electron Domain Geometry
Determine the of the Number of Electron Domains:
# e– domains
electron domain
around central atom
geometry
angles between
electron domains
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draw the Lewis structure
◆
count the e– domains around the central atom
2
linear
180°
lone pair of electrons = 1 electron domain
3
trigonal planar
120°
4
tetrahedral
109.5°
5
trigonal
bipyramidal
90° & 120°
6
octahedral
90°
single bond = 1 electron domain
multiple bond = 1 electron domain
CO2
SF4
central atom: C
2 double bonds ∴ 2 e– domains
e– domain geometry: linear
central atom: S
4 single bonds & 1 lone pair ∴ 5 e– domains
e– domain geometry: trigonal bipyramidal
BF3
XeF4
central atom: B
3 single bonds ∴ 3 e– domains
e– domain geometry: trigonal planar
central atom: Xe
4 single bonds & 2 lone pair ∴ 6 e– domains
e– domain geometry: octahedral
Molecular Geometry
Molecular Geometry
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What shape is created by the atoms that make up
a molecule?
How, geometrically, are outer atoms arranged
around the central atom?
OR
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How, geometrically, are the bonding e– domains
arranged around the central atom?
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When the e– domains around the central atom are
all bonding domains, the molecular geometry is the
same as the electron domain geometry.
Correlation Table for Molecular Geometries
linear
bent
tetrahedral
Correlation Table for Molecular Geometries
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Where do you “put” the lone pairs
of electrons in trigonal bipyramidal
geometry?
recall: the goal is to minimize
e– pair – e– pair repulsion
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a lone (nonbonding) pair of e–’s
is bigger than a shared (bonding)
pair of e–’s
∴ lone pairs of e–’s occupy
positions that are farthest away
from other electron domains
◆
in trigonal bipyramidal geometry,
equitorial positions are favored for
lone pairs of electrons
Correlation Table for Molecular Geometries
see-saw
t-shaped
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Where do you “put” the lone pairs of electrons in
octahedral geometry?
linear
square pyramidal
central atom: I
5 single bonds & 1 lone pair ∴ 6 e– domains
e– domain geometry: octahedral
molecular geometry: square pyramidal
F–I–F bond angles: 90°
central atom: Xe
2 single bonds & 3 lone pairs ∴ 5 e– domains
e– domain geometry: trigonal bipyramidal
molecular geometry: linear
F–Xe–F bond angle: 180°
square planar
central atom: Cl
3 single bonds & 2 lone pairs ∴ 5 e– domains
e– domain geometry: trigonal bipyramidal
molecular geometry: t-shaped
F–Cl–F bond angles: 90°
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Why are observed bond angles not always the same
as the predicted bond angles?
Lone pairs of electrons are bigger and occupy more
space, therefore they have the effect of compressing
bond angles slightly.
Bond Polarity vs. Molecular Polarity
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Covalent bond polarity is
determined by the difference
in electronegativity between
2 atoms.
Molecular polarity is determined
by the arrangement of bond
dipoles around the central atom.
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Why are observed bond angles not always the same
as the predicted bond angles?
Multiple bonds occupy more space than single
bonds, therefore they compress bond angles slightly.
Molecular Polarity:
Behavior of Polar Molecules in the
Presence of an Electric Field
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Bond Polarity vs. Molecular Polarity
bond dipoles are vector quantities:
have a magnitude and a direction
Molecular Polarity:
Considering Bonds Dipoles and a
Molecule’s Electron Density Map
symmetrically arranged, equal bond dipoles can
cancel each other out
CO2: 2 polar C–O bonds arranged 180° apart
∴ CO2 is nonpolar
H2O: the arrangement of the 2 polar H–O
bonds, and the 2 lone pairs of e–’s on O give
H2O a net dipole moment, μ > 0
∴ H2O is polar
Geometric Isomers and Polarity
* if dipole moment equals zero (i.e. a zero dipole moment, the
molecule is nonpolar
if dipole moment is not equal to zero (i.e. a nonzero dipole
moment), the molecule is polar
Valence Bond Theory (VBT)
Recall Covalent Bond Formation Between 2 H Atoms
Lewis theory (Ch. 9): covalent bond forms when an
electron pair is shared between 2 atoms.
Valence Bond Theory: concentration of electron
density between atoms occurs when a valence
atomic orbital on one atom overlaps with the
valence atomic orbital on the other atom.
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overlap of orbitals allows 2 e–’s to be
shared in the common space between nuclei
the greater the orbital overlap, the stronger
the bond
H2:
Consider BeF2
Lewis structure predicts (correctly) that BeF2 is
linear with 2 identical Be–F bonds.
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Consider H–Cl
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How do we describe the bonding in terms of
valence bond theory?
Which orbitals on Be and F overlap?
Why does this overlap result in linear
geometry and identical bonds?
sp Hybrid Orbitals: BeF2
Hybrid Atomic Orbitals
Atomic orbitals can mix together in a process called
hybridization:
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◆
hybridization results in formation of new
hybrid atomic orbitals:
new shape
new spatial orientations
new, degenerate energy
2p
E
e– promotion
hybridization
↑↓
2p
↑
2s
the total number of orbitals remains constant –
orbital conservation
# hybrid orbitals formed = # atomic orbitals mixed
sp2 Hybrid Orbitals: BF3
↑
sp
hybrid orbitals
shape of hybrid orbitals
sp3 Hybrid Orbitals: CH4
Valence Bond Theory Description of H2O
Hybridization Involving d orbitals:
sp3d and sp3d2 Hybrid Orbitals
XeF2 and XeF4
Prepare and Orbital Diagram for XeF2:
Orbital Diagram for XeF4:
Geometry Created by Large Lobes of Hybrid Orbitals
Sigma (σ) vs. Pi (π) Bonds:
Explaining Single, Double, and Triple Bonds
There is a specific correlation between:
◆ the number of electron domains around an atom
◆ the electron domain geometry around that atom
◆ the hybridization of the atom:
In a σ bond, e– density is concentrated symmetrically
around the internuclear axis.
◆
results from the head-to-head
overlap of orbitals
atomic orbitals (ex. s–p
or p–p overlap)
OR
hybrid orbitals (ex. sp3–sp3
or sp3d2–sp3 overlap)
Sigma (σ) vs. Pi (π) Bonds:
Explaining Single, Double, and Triple Bonds
In a π bond, e– density is concentrated in 2 parallel
regions above and below the internuclear axis.
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Sigma (σ) vs. Pi (π) Bonds:
Explaining Single, Double, and Triple Bonds
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a single bond is one sigma bond (2 e–’s)
results from the parallel
overlap (or side-to-side
overlap) of unhybridized p
orbitals
2 regions of electron density,
but still 1 bond – still just
1 pair of e–’s
for atoms to participate in π bonding, they must
have at least one unhybridized p orbital
π bonds require atoms with sp2 or sp hybridization
bond multiplicity
a double bond is one sigma bond + one pi bond (4 e–’s)
a triple bond is one sigma bond + two pi bonds (6 e–’s)
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a σ bond must form before a π bond forms
can have a σ bond without a π bond
cannot have a π bond without a σ bond
the amount of overlap in a π bond is typically less
than the amount of overlap in a σ bond
∴ a π bond is weaker than a σ bond
Bonding Description for Ethene, C2H4
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σ Bonding Description for Ethene, C2H4
each C is sp2 hybridized
each C has 1 unhybridized p orbital that can
participate in π bonding
each C–H σ bond is the result of sp2–1s overlap
the C–C σ bond is the result of sp2–sp2 overlap
the C–C π bond is the result of 2p–2p overlap
π Bonding Description for Ethene, C2H4
Bonding Description for Ethene, C2H4 :
Experimental Evidence in Support of π Bonding
Bonding Description for Acetylene, C2H2
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σ Bonding Description for Acetylene, C2H2
each C is sp hybridized
each C has 2 unhybridized p orbitals that can
participate in the formation of 2 π bonds
each C–H σ bond is the result of sp–1s overlap
the C–C σ bond is the result of sp–sp overlap
each C–C π bond is the result of 2p–2p overlap
π Bonding Description for Acetylene, C2H2
σ and π Bonding Description for Acetylene, C2H2
Using Valence Bond Theory to Interpret
Resonance Structures and Delocalized Bonding
Parallel, unhybridized p orbitals on each O are
aligned for continuous overlap and ∴ delocalization of
π electron density evenly over all 3 atoms.
HCN
CO32—
PF4—
Orbital Overlap Bonding
Description
Hybridization of Central
Atom
Polar or Nonpolar?
Bond Angles
Molecular Geometry
# Nonbonding Domains
# Bonding Domain
Template for Complete Structure, Geometry, Polarity, Hybridization and Bonding Description
Chapters 9 & 10; Lewis Structures, Resonance, VSEPR Theory, and Valence Bond Theory
Electron Domain
Geometry
Using Valence Bond Theory to Interpret
Resonance Structures and Delocalized Bonding:
Benzene, C6H6
Complete Molecular Structure-Bonding Template:
Total # of Electron
Domains at Central Atom
Chem 1711:
Lewis Structure (including
resonance if appropriate)
Each O is sp2 hybridized; the σ bonding picture is:
Total # Valence Electrons
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Recall ozone, O3 and its Lewis structure:
Chemical Formula
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Using Valence Bond Theory to Interpret
Resonance Structures and Delocalized Bonding
89!
! )(+5&,K65-$!
+6'0)/5-!Q'4R651'K5-$!
XD$!=9!
%&02!'1!2&,!)*+,(-+0.!3,*),2.4!0.*-/G!(0.Q*/!02*)!=
D9!
! -'/,56$!
@9!
◆ What is the hybridization
B9! +,+65&,K65-$!
of each numbered C?
C9! +6'0)/5-!4-5/56$!
Molecular Orbital (MO) Theory
Extension of quantum mechanics:
8F:G! 8!
>H:G! I!
LC;G! IM$A!
consider the wave functions of electrons and
N=OG! 844-R!Y5-,/(,!Q)/K!+&,)6R!S+.)^!+&6,,^!)6!*)26!,-,(+6)/!45'639$!SCP514-,!IM$<9!
a
resulting orbitals extending over a whole molecule
◆ How many σ and π
HN>G! Q)/K'/0!T!Q)/K'/0!+&,)6',3
!
bonds are in this molecule?
instead of just a single atom
! !
:;! 2,2.0&,G.0+!
#M$! D).!15/R!3'015!5/K!4'!Q)/K3!56,!'/!+&,!1)-,(2-,!4'(+26,K!Q,-).7!
molecular orbitals:
<;! 2.'3*/0+!O+0/0.!
a molecular orbital (MO) is to an molecule as an
=;! +'/,0.!
b
◆ How many sp3 hybridized
?;! 2.'3*/0+!O4.0)'G0+!
atomic orbital (AO) is to an atom
atoms are there in this molecule?
!
!
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What is the approximate
BJ;G! ,53R!
H–C–C angle (a)?
!
@;! aQ,/2!
!
◆ How many sp2 hybridized
:>5A! <!
BC5A! D! atoms are ?EFA!
there in)*G,.02,!
this molecule?H@FA! DI$D!
6<JA! B.,G'(2!2&,!)*+,(-+0.!3,*),2.4!'/!0!)*+,(-+,$!
!
◆ What are the approximate
C6BA! Q*/G'/3!R!)*+,(-+0.!3,*),2.4
bond angles (a and b)?
!
!
!
each MO holds 2 spin-paired e–’s when full
each MO has a specific, characteristic E
each MO has a specific shape and spatial orientation
!
Molecular Orbitals for H2
!"#$%&'(
!"#$%&'(
Construct a MO Diagram for H2
$,
!"#"$%&'(")*+,$-)*+ "-.$./'
!"#"$%&'(")*+,$-)*+$,"-.$./'
*0
**
MO Diagrams for 2nd Row Elements
MO Diagrams for 2nd Row Elements
MO Diagrams for 2nd Row Elements:
Molecular Orbitals Formed From p Orbitals
Interaction of s and p atomic orbitals causes the
energy levels of some MO’s to shift:
MO Diagrams for 2nd Row Elements:
N2 (after consideration of 2s–2p interaction)
Experimental Evidence That O2 Is Paramagnetic:
MO Diagrams and Interpretation
of Magnetic Properties
MO Diagrams for 2nd Row Elements
σ2p*
σ2p*
π2p
π2p*
σ2p
π2p
*
π2p
σ2p
σ2s*
σ2s*
σ2s
σ2s
An Example of a MO Diagram for a
Transition Metal Complex: