Readings for this section.
Petrucci: Section 10-7
NOTE All the coloured pictures have java-enabled rotatable models available. Click on the image to open the page containing the java applet. Make sure you have the latest version of Java installed and that it is enabled on your browser. These will work in all browsers but I find Chrome is the easiest to configure.
Experimental evidence clearly shows us that the Lewis Model of molecular bonding, while having it's merits is far from complete. Take for example the molecule Chlorofluoromethane (CH2FCl) If we draw the Lewis Dot Structure for this molecule, we get one of two possibilities:
If you're good at visualizing, you can clearly see that the two molecules depicted here are identical (Cl is red, F is green). They are superimposable. If you're not so good at visualizing, try to rotate these two models with your mouse to make them identical.
These structures seem to show that there are two different versions of this molecule, one in which the chlorine is adjacent to the fluorine and one where it is across from it. Experimental evidence shows us that there is only one molecule with the formula CH2FCl, despite there being two different ways to depict the molecule using only Lewis dot theory.
Using similar logic, we see that the molecule CHFClBr has two distinct forms and experiment shows them to have different physical properties (optical properties). So, there are clearly examples where the Lewis dot theory breaks down. There must be further theories that can explain these observations.
It turns out that the flat representations produced in the Basic Lewis structures are the problem. Molecules are not generally flat but exist in three dimensions.
Can you superimpose these? Try rotating the models in the java applet to make all the elements line up the same way in both models. You'll find it's impossible. If you line up two elements the other two will be reversed, no matter which way you try. These two molecules are not superimposible, even though they have the exact same formula. They are called enantiomers or optical isomers.
The Valence Shell Electron Pair Repulsion Theory (VSEPR), as it is traditionally called helps us to understand the 3d structure of molecules. Although we will speak often of electron pairs in this discussion, the same logic will hold true for single electrons in orbitals, and for double bonds, where one could think of the bond as consisting of two pairs of electrons. In general, the region in space occupied by the pair of electrons can be termed the domain of the electron pair. The domain is related to the orbitals we have discussed earlier (and will elaborate on later) but the two do not necessarily refer to the same thiing.
The general concept is that the pairs of electrons repel each other and wish to locate themselves as far as possible from each other about a given nucleus. Hence, for two pairs of electrons on a nucleus, the two pairs would locate themselves exactly opposite each other, forming a bond angle of exactly 180?. If three pairs exist, they will locate themselves in a plain about the nucleus at angles of 120? from each other. higher numbers of electrons form 3d arrangements as follows.
Table: Geometry and Electron Pair Arrangements. The angles given are the ideal angles for such an arrangements.
|Electron Pairs||e- pair (domain) Geometry||e- pair diagram|
We'll now go through a set of example molecules and/or ions and discuss their geometries. It is important to note that the shape of the molecules as we discuss them here is not always the same as the electron domain geometries described above.
We will consider the molecular shapes, starting with the simplest and working up to the more complicated examples. In addition, I'll mention a classification system which may be helpful in counting electron domains used herein.
For example, methane CH4 is an AX4 molecule while ammonia NH3 is an AX3E molecule. Both of these molecules have four electron domains and hence would have a tetrahedral domain geometry as listed above. However, the shape of the molecules are not the same as we will see below.
This molecule is linear. The Be does not fill its octet shell in this situation. To do so would put a large negative charge on it and a positive charge on the Chlorine atoms. This would simply not happen since Cl is so much more electronegative than Be.
|BF3||AX3||Trigonal Planar shape|
4 electron pairs
|H2O||AX2E2||bent or angular|
Note that in all these cases, the electron-domain geometry is tetrahedral. However, the molecular shape is not always so. In the case of CH4, the molecule is actually tetrahedral in shape with a perfect Tetrahedral angle of 109.5?. The next two examples have lone pairs which occupy a larger domain volume (push more on the bonding pairs) and reduce the bond angle to less than 109.5?. The last case, HF, is simply a liner diatomic molecule. There is no bond angle.
In these cases, the electron-domain geometry is always trigonal bipyramidal. However, only the first molecule is that shape with the ideal angles of 90 and 120 degrees for the axial and equatorial bonds, respectively.
In the case of SF4, there is one lone pair and four bonding pairs. The lone pair will preferentially locate itself in an equatorial position since that position has only two other pairs of electrons within 90 degrees while an axial position would have three. Thus, the molecular would be see-saw shaped or the more technically correct name, disphenoidal. The bond angles would be less than the ideal angles of 90 and 120 degrees.
ClF3 has two lone pairs and they both locate themselves in equatorial positions for the same reasons as described in the previous case. This molecule is T-shaped with bond angles of less than 90 degrees.
In all these cases, the electron-domain geometry is octahedral and in the case of SF6, so is the shape. The molecule ClF5 has one lone pair and five bonding pairs but since all positions in the octahedral geometry are equivalent, it doesn't matter which position the lone pair takes. I drew it on the bottom position here for visual effect. In the case of XeF4, the two lone pairs will locate themselves on opposite sides of the square planar molecule. In the case of the XeF4 molecule, the lone pairs will orient themselves in a square plane and the molecule will be linear in shape.