UNIT 9 OUTLINE
Molecular Geometry and Bonding Theories
Lewis Structures are useful because they tell us how the atoms in a molecule are
connected. However, Lewis Structures do not tell us anything about the shape of a
molecule. For example, let’s look at methane, CH4:
The bonds in methane are covalent bonds (carbon is sharing electrons with each
hydrogen). To share electrons, the 1s orbital on hydrogen must overlap with the empty p
orbitals on carbon. We know that the p orbitals are orthogonal (90 degrees apart from
each other). Thus from the information we have learned in chemistry so far, we would
assume that all the atoms in methane are in the same plane and each HCH bond angle is
90 degrees. This not true.
What do we know about electrons? They have the same charge so they want to be as far
apart from each other as possible. This means that the electrons participating in covalent
bonding want to be as far apart from each other as possible. If we have a central atom, C,
with four atoms attached, the bonds can be a maximum of 109.5 degrees away from each
other. So, in reality, each HCH bond in methane is 109.5. This causes methane to look
like a tetrahedron. How can this happen if the p orbitals on carbon are at 90 degree
angles? We need a new model…
TETRAHEDRON
Electronic Geometry: Tetrahedral
Molecular Geometry: Tetrahedral
When carbon is attached to 4 atoms, like methane, we combine atomic orbitals on carbon
to form 4 new orbitals that are 109.5 degrees apart. The formula is as follows:
Orbitals combined = Orbitals created
We want to create 4 new orbitals that are 109.5 apart. Fortunately, carbon has 4 orbitals
in its valence, outer, shell. They are 2s, 2p, 2p, and 2p. When we combine them we get:
s + p + p + p = 4 sp3
We get four sp3 hybrid orbitals. sp3 orbitals are 109.5 degrees from eachother. There is
one electron in each of these orbitals (carbon has 4 valence electrons). Each hybrid
orbital can now overlap with a 1s on hydrogen to form bonds.
What do these hybrid orbitals look like? Well, you can probably guess. They look 25%
like an s orbital and 75% like a p orbital:
http://www.ntu.ac.uk/cels/molecular_geometry/hybridization/Sp3_hybridization/
We can now represent methane like this:
Where the dark wedge represents hydrogen in the front and the hatched wedge
representing a hydrogen in the back.
We can describe methane’s molecular geometry and electronic geometry as tetrahedral.
What is the difference between molecular geometry and electronic geometry? Let’s look
at another example, ammonia, NH3.
Electronic Geometry: Tetrahedral
Molecular Geometry: Trigonal Pyramidal
From the Lewis structure, one would think that, like methane, ammonia would look like
this:
However, the nitrogen has four things attached to it: a lone pair, and three hydrogens.
These four pairs of electrons want to be as far apart from eachother as possible. So, the
electron pairs will take the shape of a tetrahedron and be 109.5 degrees away from each
other. Again, the atomic orbitals on the central atom will become hybridized to make 4
sp3 hybrid orbitals that point 109.5 from each other. So electronically, ammonia looks
tetrahedral. However, instead of 4 atoms attached, like methane, ammonia has three
atoms and a lone pair. So, it has a molecular geometry of trigonal pyramidal. For
molecular geometries, we don’t count lone pairs.
http://www.sparknotes.com/testprep/books/sat2/chemistry/chapter4section8.rhtml
(Side Note: Actually, because lone pairs are held closer to the nucleus than bonding
electrons, they take up more space. So, the HNH bond angles are actually slightly less
than 109.5)
Electronic Geometry: Tetrahedral
Molecular Geometry: Bent
Let’s take this a step further and look at water, H2O. If you draw a Lewis Structure of
water, you might think it looks linear:
Again, its not. The central oxygen has four pairs of electrons attached to it, so
electronically it is tetrahedral. Oxygen’s atomic orbitals are hybridized to 4 sp3 orbitals.
Each lone pair on oxygen is in an sp3 orbital and each bonding pair of electrons is shared
due to an overlap of an sp3 on oxygen and an s orbital on hydrogen.
Because the oxygen has two atoms attached, water has a bent molecular geometry:
http://www.sparknotes.com/testprep/books/sat2/chemistry/chapter4section8.rhtml
(Side Note: Actually, because lone pairs are held closer to the nucleus than bonding
electrons, they take up more space. So, the HOH bond angle is actually slightly less than
109.5)
TRIGONAL PLANAR
Electronic Geometry: Trigonal Planar
Molecular Geometry: Trigonal Planar
Let’s look at formaldehyde, H2CO. From a Lewis Structure you may think formaldehyde
looks like this:
However, carbon has three things attached to it: two hydrogens and oxygen. The farthest
apart three pairs of electrons can be around a central atom is 120 degrees. How do we
make three orbitals that point 120 degrees away from eachother? We make hybrid
orbitals by combining the atomic orbitals around carbon.
s + p + p = 3 sp2
These orbitals look like 33% s and 66% p. They point 120 degrees away from eachother.
There is one electron in each of these orbitals. However, carbon has four valence
electrons Where is the last electron?
Notice that there is an extra p orbital on carbon that we did not hybridize (there is one
electron in this orbital). Because the oxygen has three things attached to it, two lone pairs
and a carbon, its valence shell atomic orbitals are also sp2 hybridized. This means that
oxygen has an empty p orbital just like carbon does. This p orbital also has one electron
in it). The p orbital on carbon and the p orbital on oxygen are in the same plane. They
don’t directly overlap (touch), but they overlap through space:
http://wps.prenhall.com/wps/media/objects/3311/3391094/blb0906.html
Orbitals on different atoms that directly overlap by touching are called sigma (σ) bonds.
Orbitals on different atoms that overlap by being in the same plane but do not touch are
called pi (π) bonds. In this case, there is one electron in each unhybridized p orbital on
carbon and oxygen. They are shared by a pi bond.
Question: Are pi bonds weaker or stronger than sigma bonds?
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Question: What is a triple bond?
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http://www.chemtube3d.com/orbitalsformaldehyde.htm
So the double bond between oxygen and carbon consists of a sigma bond and a pi bond.
Question: Are the bond angles in formaldehyde equal? Why or why not?
________________________________________________________________________
To summarize, formaldehyde has a trigonal pyramidal molecular geometry due to three
things attached to the central carbon. The electronic geometry is the same because there
are no lone pairs on the central carbon. Let’s look at a situation where there are lone
pairs on the central atom.
Electronic Geometry: Trigonal Planar
Molecular Geometry: Bent
Lets look at nitrite, NO2-. You may think it looks like this:
However, the central nitrogen atom has three things attached: a lone pair and two
oxygens. The farthest apart three things can be around a central atom is 120 degrees.
Thus, like the previous example, the atomic orbitals around nitrogen must be sp2
hybridized and electronically it is trigonal planar. However, the molecular geometry is
bent because we don’t count lone pairs when determining molecular geometry.
Exercise: Draw the structure of nitrite.
Question: Is the ONO bond angle exactly 120 degrees?
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Question: What is the hybridization of the oxygens?
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LINEAR
Electronic Geometry: Linear
Molecular Geometry: Linear
Let’s look at carbon dioxide: CO2. You may think carbon dioxide looks like this:
You would be correct! Carbon has two things attached, two oxygens. The farthest apart
two things can be around a central atom is 180 degrees. So, we need two orbitals that
point 180 degrees away from eachother:
s + p = 2 sp
We get two sp orbitals. What do they look like? 50% s and 50% p.
http://wpscms.pearsoncmg.com/au_hss_brown_chemistry_1/57/14648/3749964.cw//3749998/index.html
We hybridized an s and a p orbital on carbon. We have 2 p orbitals left. The
hybridization of the oxygens are sp2 leaving one p on each oxygen left (can you see why
this is true?) Each unhybridized p orbital on carbon can overlap with one on each oxygen
to form two pi bonds.
http://www.thoughtyoumayask.com/picsbtqq/hybridization-of-oxygen-in-co2
The Rest of the Story
For larger molecules we are going to abandon the idea of hybrid orbitals because it gets
really complicated. To figure out the shapes of large molecules, we will focus merely on
the fact that electron lone pairs and bonding pairs like to be as far apart as possible.
As we learned in the last unit, some atoms can have more than an octet (S, P, etc). They
can have up to 12 pairs of electrons or six atoms attached. We learned about central
atoms having 4 pairs electrons. Now we add a pair.
Electronic Trigonal Bipyramidal
Molecular Geometry: Trigonal Bipyramidal
The farthest apart 5 pairs of bonding electrons can be around a central atom gives the
following shape. This is an example of PCl5:
However, what if one of these atoms were replaced with a lone pair as in SF4?
Electronic Trigonal Bipyramidal
Molecular Geometry: Seesaw
In SF4, S has one lone pair and 4 bonding pairs of electrons. Electronically it is still
trigonal bipyramidal, however, it now has a “seesaw” molecular geometry.
Question: Why is the lone pair equatorial and not axial?
______________________________
Let’s stay with the trigonal bipyramidal electronic geometry, but replace one more atom
with a lone pair.
Electronic Trigonal Bipyramidal
Molecular Geometry: T-shaped
ClF3 has a trigonal bipyramidal electronic geometry but a “T-shaped” molecular
geometry.
This next picture will give you a better idea of where the lone pairs are:
http://glossary.periodni.com/download_image.php?name=tshaped_molecular_geometry.png&source=T-shaped+molecular+geometry
Question: Why are the lone pairs equatorial and not axial?
_____________________________
Finally let’s replace another bonding pair of electrons with a lone pair.
Electronic Trigonal Bipyramidal
Molecular Geometry: Seesaw
XeF2 is electronically trigonal bipyramidal because Xe has five things attached: two
fluorines and three lone pairs. However, it is geometrically linear.
Question: Where are the lone pairs and why?
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OCTAHEDRAL
Electronic Geometry: Octahedral
Molecular Geometry: Octahedral
The farthest apart electron pairs can be around a central atom is an octahedron. SF6 is an
example.
There are six atoms attached to the sulfur atom and no lone pairs thus the electronic and
molecular geometry are octahedral. What if we replace one of the atoms with a lone pair?
Electronic Geometry: Octahedral
Molecular Geometry: Square Pyramidal
Our example for this molecular geometry is BrF5. Br has 6 pairs of electrons attached so
it has an octahedral electronic geometry. However, it only has 5 atoms attached (one pair
of electrons is a lone pair). Thus the molecular geometry is square pyramidal:
Question: Where is the lone pair?
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Electronic Geometry: Octahedral
Molecular Geometry: Square planar
Our example for this molecular geometry is XeF4. Xe has 6 pairs of electrons attached so
it has an octahedral electronic geometry. However, it only has 4 atoms attached (one pair
of electrons is a lone pair). Thus the molecular geometry is square planar:
Question: Draw the structure of XeF4. Where are the lone pairs? Why?
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Bond Strength and Length
Order of length: single > double > triple
Order of strength: single < double < triple
Resonance and Pi Bonds
Question: What is delocalization? How does the concept of resonance explain this?
(Hint: Draw all resonance structures of benzene (C6H6). Draw all the unhybridized p
orbitals. Where are they oriented in space?)
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Molecular Orbital Theory
In addition to Hybridization Theory, there is another theory of bonding, Molecular
Orbital Theory. When solving the Schrodinger Equation for the electron, you obtain
mathematical functions that describe waves that correspond to atomic orbitals: 1s, 2s,
2p… (Side note: the square of these functions describe the probability density shapes of
s, p, and d we studied earlier.) Remember this picture:
http://hyperphysics.phy-astr.gsu.edu/hbase/ewav.html
You can think of these atomic orbitals as waves. These waves can be positive in one
region and negative in another region. Do not think of these negatives and positives as
charges; the solutions to these wave equations are mathematical.
Bonding occurs by the interaction of these orbitals. We can think of the interaction as
standing waves:
http://method-behind-the-music.com/mechanics/physics
Constructive
Destructive
These orbitals can interact destructively or constructively. Constructive interaction gives
a resulting wave of higher amplitude. When waves interact destructively, they cancel.
Waves (orbitals) that interact constructively are called molecular bonding orbitals. Waves
(orbitals) that interact destructively are called molecular antibonding orbitals. Now
prepare for the weirdness: these atomic orbits combine in both ways simultaneously. The
most important thing to remember is that the number of waves that interact equals the
number or molecular orbitals they form:
Atomic orbitals in = Molecular orbitals out
Antibonding orbitals are less stable than bonding orbitals. When atomic orbitals combine
to form molecular orbitals, antibonding orbitals and bonding orbitals are formed.
Electrons always exist in the bonding orbitals first.
Example: Hydrogen gas
http://www.science.uwaterloo.ca/~cchieh/cact/c120/mo.html
Molecular orbital theory can tell you if a molecule is stable or not.
Question: Is He2 stable?
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http://www.chem.ucalgary.ca/courses/350/Carey5th/Ch02/ch2-2-3.html
Question: Fill in the following molecular orbital diagram for Oxygen gas, O2. Is oxygen
gas stable?
http://www.chem.queensu.ca/people/faculty/mombourquette/firstyrchem/molecular/orbit
als/index.htm
Notice there are π and σ symbols. σ represents orbitals that come into direct contact with
each other to form molecular orbitals. Π represents atomic orbitals that do not come into
direct contact; they are pi bonds.
The closer in energy two orbitals are, the more strongly they interact and the lower the
energy of the bonding orbital. The antibonding orbital is also correspondingly higher in
energy.
Question: Describe the interactions between these sets of orbitals:
Rules for MO Theory:
1) The most important thing to remember is that the number of atomic orbitals that
interact equals the number or molecular orbitals they form.
2) Atomic orbitals combine most effectively with other atomic orbitals of similar energy.
3) The effectiveness with which two atomic orbitals combine is proportional to their
overlap. As overlap increases, the energy of the bonding MO is lowered and the energy
of the antibonding MO is raised.
Bond Order: Tells you about the stability of a covalent bond:
Bond order = ½(# of bonding electrons - # antibonding electrons)
A bond order of 1 represents a single bond. A bond order of two represents a double
bond.
Magnetism
Paramagnetic: A molecule that has an unpaired electron is paramagnetic. They are
attracted to magnetic fields.
Diamagnetic: A molecule that has no unpaired electrons. It is weakly repelled by a
magnetic field.
Question: Is O2 diamagnetic or paramagnetic? Fill in the molecular orbital diagram above
to explain your answer.
Acids and Bases
Lewis acid:
Lewis base:
Organic chemists call Lewis acids electrophiles and Lewis bases nucleophiles.
Question: Which is the Lewis acid and which is the Lewis base?
http://en.wikipedia.org/wiki/Lewis_acids_and_bases
(Pictures lacking a reference were taken from Wikipedia)