Grade12 AP Chemistry

Tuesday May 6th

Complicating Issues with Bonding
So, we can follow three sets of rules and come up with a Lewis structure that is a pretty good 
representation of a naturally occurring molecule. Are there any situations in which Lewis structures 
fail to produce something that matches reality?  In a word... Yes!

Resonance Hybrid Structures
There are many formulas that yeild more than one structure with the same atoms joined together in 
the same way, but with a slightly different arrangement of the bonds, each of which are just as 
preferable with respect to their formal charges.

One example is the ozone molecule which, when we follow the steps yeilds two possible molecules:


At the last step of the process, we can promote a lone pair to a double bond to satisfy the stable 
octet of the central atom, but we can do so from either of the terminal oxygens to yield equally
preferable formal charges.  So, we have two structures (from the standpoint of the particular 
oxygen atoms) for the same formula.  We are not just flipping the picture around, in each case a 
different atom has the double bond.

Another example is the carbonate ion from last night's homework.  There are three oxygens from which
a lone pair could be promoted to make a double bond.  


Predicted Molecular Properties
Using the Lewis structures generated for ozone and carbonate ion, we would predict that the ozone 
should have one short bond (the double bond) and one longer bond (the single bond) and the carbonate
ion should have 1 short bond and two long bonds.  Furthermore, the carbonate ion should have a greater
bond angle between the double bond (which has a greater negative charge density) and the single bonds.  


Actual Molecule Properties
Oddly enough, when we investigate these molecules using one of methods mentioned last class, we find
that the all the bonds in each molecule are of equal length and are shorter than the predicted single 
bond and longer than the predicted double (for ozone, the bonds are 128 nm unlike the 120 pm for
a double bond and 147 pm for a single bond).  In addition, the bond angles in carbonate all are equal.

It is almost as if the molecules were rapidly transitioning back and forth between the two Lewis
structures by shifting electron pairs between bonds or lone pairs.
Lewis structures like this are called Resonance Structures.


Resonance Hybrids
In fact, we do not have any evidence that the molecule is actually resonating between the two states.
Instead, the evidence points to a stable molecule with two equal-length bonds with fractional bond orders.
The bond order of the ozone molecule is one and a half.  

In addition, the two terminal oxygen atoms seem, on investigation, to have one and one half lone pairs.
A similar situation is found for the carbonate ion, but the bond order is 1 and one third with 2 and 
two thirds lone pairs per oxygen.


The best explanation at present is that one of the bonds is being spread out between more than two
atoms.  It must be that by doing this, the charges are better balanced over the whole molecule and
the potential energy of the molecule is reduced.

Unfortunately, the Lewis model of covalent bonding does not deal with fractional bond orders, nor with
what appear to be partial lone pairs.  Obviously, we need another way of looking at covalent bonding!
Here is another problem for which an orbital model of covalent bonding should help us out.

In the mean time, when writing Lewis structures for molecules that have more than one equally
preferable configurations, we will write each of the resonance structures out with arrows between
them to indicate that the actual molecule is a configuration that is an "average" or "hybrid" of the 
individual resonance structures, at least as far as the bond orders are concerned.


Practice
Write out the structure for the acetate ion (H3CCOO1-).
Is there any possibility for resonance in this molecule?
What evidence would you use to support this claim?
Is there any possibility for resonance in a molecule of acetic acid (H3CCOOH)?


Odd Electron Species
Most molecules have an even number of electrons, or they would react with other odd electron molecules
or atoms to complete its octets.  This is one of the central concepts in the construction of Lewis
diagrams.

However, there are many, albeit very unstable (reactive), molecules that have unpaired electrons. These 
species often are produced in high energy environments or in certain chemical reactions.  Some examples
are the neutral molecules OH, NO, and CH3.  

Practice
Write out Lewis diagrams for each of formulas above.  Each one will have an unpaired electron.


Incomplete Octets
Atoms of Aluminum, Boron, and Beryllium do form molecules, but they do not have enough unpaired electrons
to make the four bonds needed for a complete octet.  They also do not have a high enough electronegativity
to make a co-ordinate bond with another atom that could supply a fourth pair of electrons.  They do, 
however, seem to be able make stable molecules in which they have less than a stable octet.

A good example is BF3
The classical structure for BF3 is the top left-hand structure in the diagram below.

One bit of supporting evidence for this structure is that BF3 can co-ordinately bond to F-
to make BF3- (bottom center structure in the diagram above).  The boron in BF3
can make a co-ordinate bond because the attraction of its electrons towards the Fluorine atoms (they are very
electronegative) gives the Boron a partial positive charge and gives it a stronger attraction for lone
pair electrons.

However, there is some evidence from other investigations that the Boron makes a double bond with one of 
the Fluorines (top middle structure).  The evidence is that the B-F bonds in BF3 are shorter 
than predicted, as if the BF3 was a resonance hybrid.  However, the formal charges for this 
structure are not sensible.  Moreover, this structure would have little tendency to form BF4-.

Practice
Calculate the formal charges for the middle BF3 structure.  Why are these formal charges not likely?

Yet another possible structure for BF3 is shown by the inonic structure in the upper right of the 
diagram above.  What sort of evidence would we look for to support an ionic or partial ionic character for
this substance?  It may even be that Boron trifluoride is a resonant hybrid of all three of the upper 
structures.  In spite of the uncertainty in the structure of Boron trifluoride, there is good evidence to 
support the existance of uncomplete octets in some Boron, Beryllium, and Aluminum compounds.

Expanded Octets
On the other side of the spectrum, there are many elements which can form stable molecules with more than 
8 electrons in their valence shell.  Some examples are:  PCl5, SF6, and even XeF6.
(yes, you read that right, Xenon hexafluoride).  The question is, where do the extra electrons fit?  An
orbital explanation might help, given that Phosphorus and Sulfur are in the third period and, thus, have
additional orbitals available in their valence shell.


Expanded valence shells have been applied to common molecules, such as the sulfate ion, to produce 
structures that break the stable octet rule, but result in lower formal charges.
The left-hand structure for sulfate is what one usually sees and what we get by following rule set number
2 from last class.  It does have very large formal charges, however.  The right-hand structure reduces 
the formal charges, at the cost of increasing the number of valence electrons for Sulfur.  As far as I
can tell, the issue is not resolved yet.  Just one of the many exciting problems to be solved by future
chemists, such as yourselves.  


Practice
Calculate the formal charges for both versions of the sulfate ion structure.

Other Issues
There are other issues for which the Lewis model of bonding cannot account.
1) O2 is paramagnetic when all of its electrons in the Lewis structure are paired.
2) Li2 and B2 exist, but Be2 does not.