Transcript electron–domain geometry

Chapter 9 – Molecular Geometry
and Bonding Theories
Homework:
11, 13, 15, 19, 20, 21, 25, 26, 31, 34,
35, 36, 39, 41, 42, 43, 44, 47, 49, 51,
54, 56, 96, 100
9.2 – The VSEPR Model
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Two balloons
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Three balloons
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Linear arrangement
Trigonal-planar arrangement
Four balloons
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Tetrahedral arrangement
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Electrons in molecules behave like balloons
A single covalent bond forms between atoms
when a pair of electrons is between the
atoms
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A bonding pair of electrons defines a region in
which the electrons are most likely to be found
between two atoms
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This area we find electrons is called an electron domain
A nonbinding pair (or lone pair) defines an
electron domain located around one atom
Example
• Four electron domains here
•In general, each nonbinding pair, single bond or multiple
bond produces an electron domain around the central atom
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Because electron domains are
negatively charged, they repel each
other.
The best arrangement of a given
number of electron domains is the one
that minimizes the repulsions between
them.
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This is the basic idea behind the VSEPR
model.
Similar to Balloons?
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You bet!
Two domains makes linear arrangement
Three domains makes trigonal-planar
arrangement
Four domains makes tetrahedral
arrangement
pg. 349
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The arrangement of electron domains
about the central atom is called its
electron–domain geometry.
In contrast, the molecular geometry is
the arrangement of only the atoms in a
molecule or ion
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So any non-bonding pairs are not a part of
the molecular geometry
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The VSEPR model predicts electron-domain
geometry
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From this and knowing how many domains are
due to nonbinding pairs, we can predict the
molecular geometry
When all the electron domains in a molecule
come from bonds, the molecular geometry is
the same as the electron-domain geometry
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But if one or more domains comes from lone pairs,
we must ignore those domains for molecular
shape
pg. 351
Example
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NH3
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Already done this.
4 electron domains around central atom
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So electron-domain geometry is tetrahedral
We know 1 of those domains comes from
lone pairs
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So the molecular geometry of NH3 is trigonal
pyramidal
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Tetrahedral with one less end, see pg. 347
Steps using VSEPR model to
predict shape of molecules
Draw Lewis structure
1.
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Count number of electron domains around
central atom
Determine electron-domain geometry
2.
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Use table 9.1, 9.2 or 9.3
Use the arrangement of the bonded atoms
to determine the molecular geometry
3.
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Use table 9.2 or 9.3
Example
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1.
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CO2
Draw Lewis Structure
How many electron domains around the
central atom are there?
What is the electron-domain geometry for
this?
2.
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Linear
What molecular geometry is possible?
3.
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Linear
Effect of Nonbonding Electrons and
Multiple Bonds on Bond Angles
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We refine the VSEPR model to predict
and explain slight variances from the
ideal bond angles
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Methane (CH4), ammonia (NH3) and water
(H2O) all have tetrahedral electron-domain
geometries
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But their bond angles are a little different
CH4 = 109.5º, NH3= 107º and H2O = 104.5º
Differences based around which type of
electron pairs make up the electron domains
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Bond angles decrease as the # of
nonbonding electron pairs increase.
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Bonding pair of electrons attracted by both
nuclei of the bonded atoms
Lone pair of electrons attracted primarily by
one nucleus
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Because lone pair has less nuclear attraction,
it’s domain becomes more spread out
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So electron domain for lone pairs exert more
repulsive force on adjacent electron domains
This compresses (lessens) the bond angles
Since H2O had the most lone pairs, it gets the
shortest bond angles
Multiple Bonds an Bond
Angles
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Multiple bonds have a higher electroncharge density than single bonds
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Also creates larger electron domains
So electron domains for multiple bonds
exert a greater repulsive force on adjacent
electron domains than single bonds do
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So multiple bonds (double or triple) will
decrease the bond angles too
Phosgene (Cl2CO)
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Central atom has three
electron domains
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3 single bonds
Trigonal planar geometry
Double bond acts like a
lone pair, reducing the
Cl-C-Cl bond angle
How Do These all Compare?
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In terms of volume occupied by electron
pairs
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In other words, who compresses the most?
Lone pair > triple bonds > double bonds >
single bonds
Molecules with Expanded
Valence Shells
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So far we have assumed the molecules
have no more than an octet of electrons
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But the most common exception to the
octet rule is a central atom having greater
than 8 valence electrons
So we need to deal with molecules with 5
or 6 electron domains
pg. 354
Example
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Use the VSEPR model to predict the electron
and molecular geometry of ClF3
Step 1: Lewis structure
How many electron domains around central
atom?
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5
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5 electron domains
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How many bonding domains?
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3
How many non-binding domains?
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Gives us an electron geometry of trigonal
bipyramidal
2
So its molecular geometry is
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T-shaped
Shapes of Larger Molecules
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The VSEPR model can be extended to
more complex molecules than we’ve
been dealing with.
Consider acetic acid
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CH3COOH
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Acetic acid has 3 interior atoms
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Carbon, and each oxygen
We can use VSEPR to look at each
central atom individually
9.3 – Molecular Shape and
Molecular Polarity
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Remember that bond polarity measures
how equally the electrons in a bond are
shared between the two atoms
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Higher bond polarity = less equal sharing
Higher electronegativity difference = higher
bond polarity
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The dipole moment depends on both
the polarities of the bonds and the
geometry of the molecule
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Last chapter we focused just on the
polarity effect on the dipole moment
For every bond in the molecule, we can
look at the bond dipole
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The dipole moment that is due ONLY to the
two atoms in the bond
Example
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CO2
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O=C=O
Each C=O bond is polar (O is more
electronegative than C)
Since we have two O=C bonds, the bonds
are identical
We end up with high electron density
around the O, and low electron density in
the middle
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Bond dipoles and dipole moments are vectors
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The overall dipole moment is the sum of the bond dipoles
that make it up
But, must consider both the amount of the dipole, and
the direction of the dipole
We have two identical C=O bonds, so the amount of
the dipoles are the same
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But the DIRECTION of the dipoles are opposite
This causes the individual bond dipoles to cancel each other
out
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So the geometry of CO2 indicates that it is a NONPOLAR
molecule, even though it contains polar bonds.
Bond Dipole Activity
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Bond Dipole Activity
Steps to Determine Molecular
Polarity
1.
2.
3.
Draw Lewis structure
Determine molecular geometry
Look at effects of electronegativity
differences
9.4 – Covalent Bonding and
Orbital Overlap
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The VSEPR gives as a method to predict the
shape of molecules
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Does not explain WHY the bonds exist between
atoms
A mixture of Lewis’ notion of electron-pair
bonds and atomic orbitals leads to a model of
chemical bonding
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This mixture of views is called the valence-bond
theory
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In Lewis theory, covalent bonding occurs
when atoms share electrons
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The sharing concentrates electron density
between the two nuclei involved
In valence-bond theory, the build-up of
electron density between the nuclei is thought
of as occurring when
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a valence atomic orbital of one atom merges with
a valence atomic orbital of another atom
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This merger of orbitals
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Means that they share a region of space
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The overlap of orbitals allows two electrons
of opposite spin to share the common
space between the nuclei
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Called overlap
Forming an atomic bond
See figure 9.14 on pg. 360
Distance
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There is always an optimum distance
between the two bonded nuclei in a
covalent bond
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Too close = too much repulsion between
the nuclei
Too far = not much overlap, not a strong
bond
9.5 – Hybrid Orbitals